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Chairman, Drpmtmrnl of j'hyuftil Science 
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. For success and satisf action in modern living, tin* citi/en 
must understand and appreciate the scientific aspects of his 
environment and the role ol science in the development of 
civili/ation. lie must hase his everydav thinking and action 
on the best information available with a full reali/ation of 
why the methods of science are superior to other methods of 
obtaining information. Finally, the citi/en must be ready to 
use scientific knowledge lor the <M)od of all. M \v \\'ORI,|) or 
CIII.MISTKY is designed to help students achieve these objec- 
tives. It is a textbook for yoimjj; people who are learning to be 

This 1 ( ) r ).") edition of MAV \VOKI.D 01- CIIMMISTRY lias been 
brought abreast of recent chemical developments and dis- 
coveries. Striking advances have been made in nuclear energy, 
petrochemistry, metallurgy, textiles, and plastics. The chapters 
dealing with these subjects have been tlionni^lHy revised to 
take these changes into account. 

The basic plan ol the book, which has been so widely 
accepted, lias not been altered. Hut within that framework 
many changes have been made. Diagrams have been com- 
pletely redrawn and enlarged. New illustrations have been 
provided. The index lias been expanded to make the text an 
even more effective source book. Despite the addition of new 
materials, however, the hook as a whole lias been shortened. 
Perhaps most important of all. the chapters dealing with 
the basic theory and mathematics of chemistry have been 
rewritten. Formulas, valence, equations, and problems have 
been iven a new and expanded treatment. The chapters pre- 
senting the electron theory, ioni/ation. and the study of teases 
have been rewritten to achieve a maximum of clarity. While 
the concepts of chemistry themselves cannot be simplified. 


thry can he made r.isiri to understand h\ hcttcT arrangrmrnt 
and tfn.ilci clanlx of laniiua^c' Glarit) of wilting has been 
.L m.ijoi 140.1! of this substantial levisinn of NKW WORLD OF 


'I he bock has been \\iitten \\ith the 1 puipose of getting idea* 
acrms to the \tudcnL With the exception of ncccssaix teehnical 
trims, tin \o(,ibulai\ consists of familial words. The style is 
duett, tin sentences and paia<;iaphs are short. Much of the 
illiistiatm mateiial has been draxxn fiom the student's own 
enviionmc nt (lonsumii aspects of che-mistix and the them- 
isti\ of common things aie stiessed at the same time that 
ic(|iiiienie.its of model n couises of stud\ and examining boards 
aie iulh met. 

Select! d matenals from the lnstor\ of c hemistr\ arouse 
inteiest .ind enable the student to sec how fundamental scien- 
tific ideas ha\e developed and expanded. Thc\ revc\il that 
rliemic .il thioiies ,md principles arc* descriptive generali/ations 
of ni. ill's oinv ing i xpiMieni i in tr\mi> to undei stand his world. 
FinalK, the 1 hi^tonc materials hel]) studrnts scu- hem scientific 
discoxm has aflerte-d modem civili/ation and hem the 1 needs 
of sot iet\ c'onstantK stimulate* s( lentific rest-. in h 

The* authoi \\islics to expiess his thanks to the mail} 
teachcTs. diemisK and industrial, e'diu and govc-rn- 
menttil or^.un/ations that have 1 generousU made suggc-stions, 
chicked the acciuacA of mateiials. and provided illustiations 
and data foi this nc\\ edition. He aNo \\ishes to ac'kncm ledge* 
his indebtedness to his son. I)i Lionel F Jaffe-, and to Lrc 
I)eihton and John H \Villi.iniscm of Silver Buidett (Company 
\\liosi- i-aieliil plaiinui!*. sound judgmc^nl, and nrvcr-f ailing 
rntliiisia^ni h.ixe made this c-oopeiative- efTort e-xtie-me^h pleas- 
ant and, it is hoped, eflrctive 

Ft nun id Jaffe 


1. MATTER and Its Changes I 

What the \\oild is made of .uul ho\\ M icntilu methods expand 
our undci^tandmg. 

2. OXYGEN: Earth's Most Abundant Element 24 

Distovei), piopeities, and uses of the gas \\lmli is the bicath of 

3. HYDROGEN: Lightest of the Elements 44 

The piimat\ stnlT of \\hkh tlie entne uni\eise is l)inlt. 

4. WATER: Most Common Liquid Compound 60 

The thiillmg discoveiy of its composition, its ptopeitics, and uses. 

5. ATOMS: Bricks of the Universe 75 

\ Quaker sc hoolteat her explains t\\o laus of <hemisti\ b\ means 
of his atoms 

6. FORMULAS: The Chemist's Abbreviations 85 

Rules on the untin^ of chemical foimulas. 

7. ATMOSPHERE: The Ocean of Air 97 

I low the components of the air \\ere discovered and utili/ed. 

8/ EQUATIONS: Shorthand of Chemistry 114 

Rules and aids in the halancing of chemical equations 

9. MATHEMATICS of Chemistry 124 

Meaning of atomic wriijhts and how some pioblems 
aie solved. 

10. CHLORINE and the Halogen Family 137 

A closelv related group of elements fluorine, c hloi me, bromine, 
and iodine. 

I L ELECTRONS and Other Particles 154 

Origins, development, and chemical usefulness of the latest theory 
of the structure of the atom. 


12. NUCLEAR ENERGY at Last! 176 

M.m hn.ilU p<netiates the < enter ot the atom and produces nu< lear 

13. ACIDS: Hydrochloric Acid, a Typical Acid 195 

PiopeitKs .iiid g<nei,d method of piep.ii mj a< ids UK hiding IIF 

14. BASES: Sodium Hydroxide, a Typical Base 206 

Nrutr ah/at ion of bases b\ .Kids, and salt formation 

15. SOLUTIONS: Water, the Universal Solvent 217 

Kinds of solutions, distilled and hid:\ \\atei 

16. IONS and Dissociation 231 

\\n\\ disputation explains stieimth ol acids, ludiol\Ms, electrolysis, 

17. AMMONIA and Reversible Reactions 250 

I lie lommeiu.d s\nthesis of aninioma a milestone in <hemistr\ 

18. NITRIC ACID and Nitrogen Compounds 262 

1 he (omiiKKial s\nthesis of nitiit .Kid, and mttogeii fixation 

19. MOLECULES: Avogadro's Hypothesis 276 

Imliidmu applications lo the solution of \\ni;lit-\oltinir and 
sliai^hl-\ olume pioblenis 

20. SULFUR and Hydrogen Sulfide 289 

I hen piodmtion, piopeities, and uses, .IK hiding the manuf.ieture 
of mat( lies 

21. SULFURIC ACID: The Fundamental Acid 304 

Its ( piepaiation, piopeities, and mam uses. 

22. ALLOTROPIC CARBON: Key Source of Energy 320 

Piopeities and uses \\ith spei attention to coal 

23. CARBON DIOXIDE: Gas of Life and Decay 339 

Including caibonic aeid, its salts, and baking poudeis. 

24. CARBON MONOXIDE and Other Gaseous Fuels 352 

Including water gas, producer gas. coal gas, natuial gas, acctvlenr 

25. METALS and Their Chemical Activity 369 

Including the piepaiation, piopcities, uses, and detettion In the 
spectroseope of sodium, potassium, .ind lithium. 


26. ALUMINUM: Most Common ol Light Metals 387 

I IK llldllli! MIMIC <>t Its compounds, .Hid j|s {\\c Ill.t 

27. IRON and Steel 404 

Intituling .1 ulmipse of the ne\\ linn/mis in the steel mdiis(i\. 

28. COPPER: Nerves of the Machine Age 423 

Its met. il hit <j,\ and uses, iiu hiding coppei siill.itr. 

29. OTHER METALS and Their Uses 436 

Including the nxe of the M>-<.ilh(i iaie metals. 

30. FERTILIZERS and Salts of Sodium 464 

IIK lulling the solution ol oin pot.ish piohlcin. 

31. CALCIUM: Its Common Compounds 480 

u supc iphiisph itcs .UK! \\.itrr 

32. IRON: Some Special Compounds 496 

I IK ludinu then uses in inks .ind l)luc|)iint ]).per. 

33. GLASS and Some Silicon Compounds 505 

IIH liiciinu ,dso some hoTon ( on i pounds MM h .is borax. 

34. HYDROCARBONS and Their Derivatives 522 

n^ the s|)C( t.n nl.ii iis<- of petioi heinistix .1 new 

35. ALCOHOL and Other Organic Compounds 547 

Anothei i*limp\r of th<- \\cnlcl of the OII^.IIIM < hnnist. 

36. FOODS and Chemotherapy 566 

Chemistry .it \\oik in the ser\ue of m,m\ he.dth 

37. FIBERS AND PLASTICS: Textiles, Paper, and Dyeing 591 

Including the ne\\ \\oild of in.iii-ni.idt 

38. COLLOIDS: The Colloidal State of Matter 610 

The rhcnustiy of tiny paitic Irs drops. fiLiincnts, L^i.uns, and films 

39. LIGHT: Its Chemical Effects 623 

Including the principles of sunple photnt^i.tphy 


Simplest formula and true formula di trniimmc; atomic s\eit^ht ex- 
perimentally, temperature s<ale <on\ersjon, use of Hoyle's Jaw and 
(ihaile^ la\>. 



. . . We must trust in nothing but 
facts. These are presented to us by 
nature and cannot deceive. We ought 
in every instance to submit our 
reasoning to the test of experiment. 
It is especially necessary to guard 
against the extravagances of imagina- 
tion which incline to step beyond 
the bounds of truth. Antoine Lau- 
rent Lavoisier, 1743-1794 

Progress through scientific knowledge. A play, presented in New 
York City, is seen and heard instantly and simultaneously in millions 
of homes throughout the nation, as clearly as if its vast audience 
were actually seated in the studio. An airplane takes off from an 
airfield in California and within a few minutes is screaming through 
the stratosphere 60 thousand feet above the earth at more than 700 
miles per hour. In an Alabama farmhouse, a country doctor diag- 
noses his patient's illness as pneumonia, and assures the family that 
it is quickly curable by the administration of certain drugs. In an 
isolated group of islands in the South Pacific, a bomb explodes with 
a force equal to hundreds of thousands of tons of dynamite. In Iowa, 
a housewife cooks a salt-water fish caught by her husband during the 
family vacation six months before and two thousand miles away. 

A generation ago, any of these occurrences would have been con- 
sidered miraculous. Yet today they arouse no unusual public excite- 
ment, for they are no longer extraordinary events. To us they are 
familiar happenings, representative of what we call the progress of 

Progress has been taking place in some measure ever since man 
appeared on earth. But within the past two or three centuries, 



civilization has progressed much more rapidly than in all the previ- 
ous thousands of years of man's history, with the most radical changes 
occurring within the past few decades. In the years ahead, civilization 
will continue to progress. More of the "incurable" diseases will be 
conquered; methods of transportation and communication will be 
further improved; the force which gives the atomic and hydrogen 
bombs their great destructive power will be harnessed for man's 
welfare; our everyday lives will be made safer and more com- 

What enables us to perform acts today which were unheard of 
just a few short years ago? How may we be certain that the future 
will bring forth new wonders? The answer to both of these questions 
lies in the knowledge we now possess concerning the nature and 
behavior of those things which make up the entire world and all its 
living creatures. In short, the answer lies in our scientific knowledge. 

Scientific knowledge is not something which has been created 
recently. It has been gathered by many men, known and unknown, 
during all of man's centuries on earth. Just as a snowball rolling 
downhill is at first small and slow moving, but in time increases both 
its si/e and speed, so it has been with scientific knowledge. In his 
early days on earth, man knew a limited number of scientific facts, 
increasing the number with painful slowness as the centuries went by. 

But gradually, almost imperceptibly, the number of facts became 
larger and their rate of discovery quickened. Within the past few 
hundred years, particularly during the twentieth century, they have 
been acquired at a prodigious rate. New scientific advances are an- 
nounced almost daily. A rolling snowball must finally come to rest 
at the foot of a hill, but there is no indication that scientific knowl- 
edge will ever cease to grow. Nor does there appear to be any limit 
to the progress civilization can make through the application of this 

Linde Air Products Company 

(left) The highly-trained 
professional chemist makes 
many important contribu- 
tions to civilization. 

(right) The lessons of the 
student chemical labora- 
tory find many useful ap- 
plications in daily life. 


Chemistry has helped make a new world. No single science is 
entirely responsible for our modern civilization. But chemistry ranks 
high among those which are considered the most important because 
it is a basic science^ essential to virtually every scientific study regard- 
less of its nature. In the manufacture of every product of our great 
industrial civilization, chemistry plays ajvital role. 

The work of the chemist affects all of us the physician, the 
dentist, the engineer, the public health employee, the soldier, the 
gas station attendant, the clerk, the housewife, the worker of factory, 
jninc, or farm, the schoolboy and the schoolgirl. 

While the achievements of modern chemists seem at times to be 
the result of some kind of magic, nothing could be further from the 
truth. Chemical ^magic'^is the result of years of study, hard work. 
and struggle; of burning the midnight oil when it seems a problem 
has no solution; and, finally, of a thrilling moment when everything 
hinges on one more experiment and the world spins 'round in a test 
tube or flask. Nor are all experiments rewarded with striking suc- 
cess. Sometimes they fail and the chemist must begin, wearily, but 
with determination, to retrace his weeks or months ol work in an 
attempt to discover where he went astray. 

Chemistry is an intensely interesting and rapidly changing field 
and, like all sciences, its roots extend far back into history. Today's 
chemist owes a great debt to the men who, over many centuries, 
added to the knowledge he now uses in performing his "miracles." 
The lives of those men, how they accumulated information and how 
that information is applied in modern chemistry is a thrilling story, 
filled with suspense and drama. It is the story told in NKW WORLD 01- 


Although we live in the "scientific age," we still see, everywhere 
in the world, superstitions and prejudices, hunger and disease. To 
correct these conditions, we must use in the outside world what we 
learn in the classroom. We must base our actions upon scientific 
knowledge, always bearing in mind, however, that we know little 
with absolute finality. 

Board of Education, City of New York 


We must also remember that science is a two-edged sword which 
may be used either to serve man or to destroy both him and his 
^ivbrks. Although the spirit of science is essentially democratic and 
constructive, we alone can prevent its becoming an oppressively 
tyrannical and horribly destructive weapon. Not many of us may 
ever become professional scientists, the men and women who work 
in the great laboratories of industry, government, and education. 
But we can all become scientists in an even broader sense. 

We can act on the most dependable information available, using 
the searching light of science to wipe out prejudices, half truths, 
and incorrect beliefs. The methods and knowledge of science are in 
our hands. We can live more fully, more satisfyingly, more com- 
fortably, more humanely, and more intelligently only by using them. 
Look forward to the future, with the assurance that, through the 
j^fforts of each of us, science amafld w ill serve mahkTncT 

What is the world made of? For thousands of years, men have 
tried to find out what the material things of the world are made of. 
Socrates (sok'rd-tez) , a learned Greek teacher who lived in the fifth 
century B.C., believed that he might discover the answer to this ques- 
tion simply by thinking about it. One other Greek teacher is said 
to have put out his eyes so that his thinking might not be disturbed 
or influenced by what he could see about him. He wanted to give 
himself over completely to undisturbed thinking, or contemplation. 

To be sure these were extreme cases. Not all the ancient teachers 
relied wholly on pure thought as they tried to find out what the 
world is made of. Some of them, of course, were influenced by what 
they could observe. For example, Thales (tha'lez) , who was born in 
Miletus in Asia Minor in 624 B.C., noticed that water nourishes crops 
and that it is found in large amounts in the bodies of men and other 
living things. Hence he speculated, or guessed, that water was the 
basic substance from which all material things were made. 

The speculations of Thales were more important to him than his 
observations. It is said that one day, as he was walking along ab- 
sorbed in deep thought and looking up at the sky, he fell into a well. 
Thratte, a housemaid, saw the accident and laughingly said: "In his 
zeal for things in the sky, he does not see what is at his feet." No 
doubt, most of you know how easy it is to become so absorbed in 
trying to solve a problem that you become "blind" to common, every- 
day things and may even be called "absent-minded." In a similar 
manner, it is possible to become so engrossed in thinking about a 
problem that simple methods of solving it do not occur to the 
thinker. Some of the ancient teachers were victims of this "blindness." 

Careful experiments and 
constant observation are 
essential to the chemist's 
study of matter. 

Many of Thales' conclusions concerning what the world is made 
of were wrong. He and other scholars like him depended too much 
upon what is now called abstract thought, and altogether too little 
upon careful experimentation and observation. II they had depended 
less upon abstract thought and more upon observation and experi- 
mentation, accurate information would have accumulated much 
more rapidly. However, even though many of their conclusions were 
wrong, their work was of value. They were thinking about the things 
that surrounded them, which was more than was being done by most 
of the people who lived at that time. 

Of course, in answering any question, careful thinking is neces- 
sary, but it must be thinking based_iipon act 'uralc Jacls, which can 
comc^only from careful observation. Many of the ancient teachers 
were fine thinkers. They began with what they knew, or thought 
they knew, and reasoned logically to a conclusion. Many of their 
conclusions were false, however, because they based their thinking 
on inaccurate or incomplete information. 

Ours is a complex world. This age-old search to find out what the 
world is made of is not over. Far from it. You, living 2500 years after 
Thales, know that water is not the basic substance from which the 
world is made. Yet, if you were asked to name all the different ma- 
terials that make up the world, you would doubtless think you had 
been given an impossible task. Imagine listing all the materials that 
make up the rocks and soil beneath your feet, the vast expanse of 
water that covers three-quarters of the earth, the multitude of liv- 
ing things, the deep envelope of air that surrounds our globe, and 
the billions of stars, some so far from the earth that their distances 
stagger our imagination! 

At extremely low tempera- 
tures, a gas may become 
liquid or even solid. This 
flask contains liquid helium 
and solid air. Because the 
frozen air at a temperature 
of 340 F. is much "hot- 
ter" than the liquid helium 
at -452 F., it is causing 
the helium to boil and over- 
flow the flask. 

Westinghouse Electric Corporation 

Matter exists in three states. Even though you could not name all 
the substances that make up our world, you would at least know that 
some of them exist as solids, others as liquids, and still others as 
gases. In fact, probably you know that mailer, which is anything that 
has weight and takes up space, occurs in one of three conditions 
solid, liquid, or gaseous. In which of these conditions, or slates of 
matter, any substance exists depends partly on the nature of the sub- 
stance itself, partly on its temperature, and partly on the pressure. 

Many substances exist in all three states, depending on the tem- 
perature. For example, water (a liquid) may be changed to ice (a 
solid) by cooling, or to water vapor (a gas) by heating. Changes 
from one state of matter to another by heating or cooling are very 
common. Iron, which we know as a hard gray solid, is melted in 
foundries and changed to a shimmering, silvery liquid. If its tem- 
perature is raised high enough, gaseous iron vapor is boiled oft. 
While iron vapor is something with which most of us are not fa- 
miliar, astronomers know that iron exists normally in the gaseous 
state in certain extremely hot stars. 

What is a physical change? When water changes to steam or to 
ice, only its form has been changed. Steam is a form of water; ice is 
also a form of water. When a piece of limestone is pulverized, only 
its form has been changed. The small pieces are still limestone, al- 
though their form differs from that of the huge chunks in which they 
came from the quarry. These changes in the water and in the lime- 
stone have been in form only. Neither the water nor the limestone 


was changed into another substance. A change in which the original 
substance does not change into one or more other substances is a 
physical change. 

Break a piece of wood in two. Heat a piece of iron wire in a vac- 
uum by passing an electric current through it. The heat produced 
causes the wire to glow, but, after the current is shut off, the wire 
returns to its original condition. In each case, what kind of change 
has occurred? 

How does a chemical change differ from a physical change? You 
know that when a piece of paper burns, it is completely and radi- 
cally changed. The hot gases that are given off and the ash that is 
left behind do not in any way resemble the original substance. When 
gasoline is burned in an engine, the resulting substances are entirely 
different from the liquid gasoline. Animal tissue is totally different 
from the vegetable substances from which it is made. The dull tar- 
nish on silverware differs completely from the gleaming silver. 

In these changes more than the form of the original substance 
was altered. In each case the composition of the original substance 
changed. Some form of energy, usually heat, was either liberated 
or absorbed. A change in which the original substance disappears 
(changes] and new substances are formed is a chemical change. 

Place a small quantity of sugar in a beaker or other glass vessel. 
On the sugar, pour a small quantity of sulfunc acid (a liquid which 
is discussed in detail in Chapter 13) . The white sugar changes to a 
black spongy substance, which cannot be dissolved, or absorbed, in 
water. What kind of change has occurred? 

Chemistry is the science that deals with the composition of matter 
and with the many chemical changes which matter undergoes. 

How is a substance identified? Telling one substance from an- 
other is called identifying a substance. When you try to identify a 
substance or to find out whether a substance has undergone a chemi- 
cal or physical change, you need to know its characteristics what 
the substance is like, and how it acts with other substances. 

We find out what a substance is like by asking such questions as: 
What is its color? its odor? its taste? Is it a solid? a liquid? a gas? 
At what temperature does it boil? At what temperature does it 
freeze? How hard is it? Does it conduct electricity? The answers to 
these questions are characteristics that enable us to describe and 
identify a substance. These characteristics of a substance are called 
its physical properties. 

We find out how a substance acts with other substances by placing 
it in contact with these substances and observing what occurs. We 


ask also how sunlight, electricity, and heat affect it. Our observations 
and the answers to these and other questions give additional charac- 
teristics that enable us to describe and identify a substance. These 
characteristics are called the chemical properties of that substance. 

Since two different substances never have exactly the same physical 
and chemical properties, any substance may be identified by deter- 
mining these properties. 

The ancients believed the world made of "four elements." But 
to return to our original question: What is the world made of? 
Guesses and speculations would be useless in attempting to answer 
this question. Because the ancients depended chiefly upon these pro- 
cedures and also upon inaccurate and uncontrolled experimentation 
and observation, they made little progress in answering the question. 

After Thales had suggested water, another man proposed that air 
might be another of the basic substances from which all matter was 
made. Fire, too, was suggested and later earth. Pythagoras (pi-thag'6- 
ras) , an ancient Greek thinker and mathematician who lived about 
600 B.C., is thought to have been the first European to express the 
idea that all matter was composed of these "four elements." 

These conclusions seemed to be proved by the observations of the 
early investigators. When a stick of green wood was burned, they 
saw that lire was produced, water was forced out and boiled off at 
the ends of the stick, a smoky vapor (air) was given off, and an ash 
(earth) was left behind. They concluded, therefore, that all matter 
was made up of different amounts of two or more of these four basic, 
or elementary, substances. 

The Greek thinkers, however, made a serious mistake. They failed 
to make enough observations of different substances. They did not 
make enough experiments. Consequently, their conclusions were 
wrong. Strangely enough, the idea that all matter is composed of 
"four elements" (earth, air, fire, and water) persisted until the 
eighteenth century and was considered correct by many otherwise 
well-informed persons. Even today speakers and writers often refer 
to the violent actions of air and water as "the fury of the elements." 

The universe is made of an even 100 chemical elements. Scien- 
tists now consider that the mountains, the oceans, the air, all living 
things, and even the stars and the rest of the universe are composed 
of simple natural substances that cannot be broken down, or de- 
composed, into simpler substances by the ordinary types of chemical 
change. A substance that cannot be broken down, or decomposed, 
into a simpler substance by the ordinary types of chemical change is 
an element. 

These natural elements do not all occur on earth in equal 
amounts. Taken together, 20 of them make up 99.5 percent of the 
weight of the crust of the earth. All the other elements comprise 
only O.f> percent of its weight. In all. there are 100 elements. Eight 
of these elements have been produced in laboratories by scientists. 
Probably the elements with which you are most familiar are gold, 
silver, iron, copper, nickel, tin, aluminum, sulfur, oxygen, carbon, 
nitrogen, and hydrogen. A list of all the elements is given on page 

What is a compound? Most of the substances we see, such as sand, 
chalk, cotton, table salt, and water, are not elements. Rather, each 
is composed of two or more elements so combined that (1) only 
chemical action can tear them apart, and (2) the elements of which 
each substance is composed can no longer be identified by their 
original individual properties. A substance composed of two or more 
elements .\o combined that the elements can no longer be identified 
by their original individual properties is a compound. The elements 
of which a compound is composed are said to be chemically com- 
bined, or themically united. 

Marble, for example, is a compound made up of three elements, 
carbon, calcium, and oxygen, chemically combined. The properties 
of a compound, such as color, odor, taste, form, and ability to dis- 
solve in water, are nearly always distinctly different from the prop 
erties of the elements of which it is composed. For example, pure 
cane sugar, a sweet, white, crystalline solid which dissolves in water, 
is completely different from any and all of the three elements of 
which it is composed. 

How does a mixture differ from a compound? In a compound 
the elements must be chemically united. But there are other kinds 
of substances made up of two or more elements or compounds. Al- 

or mixed, each of the original substances can still be identified by 
its original individual properties. Hence, the substances are not 
chemically united. 

A pinch of salt and a pinch of white sand stirred together make an 
excellent example <>!' one of these substances. The salt can be iden- 
tified by its characteristic taste and the sand by its gritty feel on the 
tongue and teeth. A substance composed of two or more elements or 
compounds that still retain their individual properties after they 
have been thoroughly mixed is a mixture. Some 1 of the most useful 
substances in the world, such as soil, air, petroleum, and milk and 
many other foods, are mixtures. 

The properties of a mixture are the same as the properties of the 
elements or compounds that compose it. A handful of iron powder 
mixed with a handful of powdered sulfur makes a mixture that re- 
sembles both the black iron and the yellow sulfur. If a magnet is 
passed through it, the iron clings to the magnet. If a liquid (ailed 
carbon disulfide is added to the mixture, the sulfur is dissolved. But 
if the mixture of sulfur and iron is heated, these two elements com- 
bine, forming a compound known as iron sulfide. Iron sulfide does 
not look like either sulfur or iron. It is not magnetic and does not 
dissolve in carbon disulfide. The properties of this compound do not 
resemble those of either sulfur or iron. 

Substances in certain mixtures may be separated mechanically. A 
mixture of salt and sand may be separated by adding water. The 
salt dissolves and the sand settles to the bottom. A mixture of iron 
and sulfur may be separated by passing a magnet through it or by 
adding carbon disulfide which dissolves the sulfur. 

The phlogiston theory, an erroneous explanation ol burning. 
One of man's greatest early achievements was the discovery of the 
use of fire. So strange did fire appear that for a long time men wor- 
shiped it. They considered it the force responsible for all creation. 
They pondered over its mystery and made many attempts to explain 
it. Karly alchemists thought that fire was the result of some vague 
"sulfur" which burnable substances contain. But later alchemists 
felt the need for a better explanation an explanation which took 
into account more of the facts that had been observed in burning 
^ many different substances. A statement that takes into account and 
attempts to explain observed tacts is known as a theory. 


Over many centuries, alchemists, the forerunners of modern chemists, worked in 
vain with their crude equipment to find the secrets of prolonging life and of 
making gold from base metal. 

About 300 years ago Becher (bek'er) , a German scientist, ad- 
vanced the theory that all burnable substances contain phlogiston 
(flo-jis't6n) , or "fire stuff." He said that when a substance burned, 
phlogiston left it in the form of (lame. Becher thought that the ash 
formed when a substance burned was the substance in in us its phlo- 
giston. According to his theory, substances that burn readily, leaving 
little ash. contain a great deal of phlogiston, while substances that 
burn with difficulty and leave much ash contain little. The phlo- 
giston theory was the first great theory in chemistry. 

The phlogiston theory seemed correct to the alchemists because of 
certain observations they had made. A rising candle flame seems to 
tug at the wick. To the alchemists this suggested that phlogiston was 
escaping from the binning candle. \Vhen a small amount of pow- 
dered lead is heated in an iron spoon, it melts, burns, and forms a 
yellow powder. According to the phlogiston theory, this yellow pow- 
der is lead ash, or lead minus its phlogiston. Now if some way could 
be found to add phlogiston to this lead ash, lead should be produced 
again. Perhaps this could be done by heating the lead ash on some 
substance that contains a lot of phlogiston, such as charcoal. The 
charcoal might give up some of its phlogiston to the lead ash. When 
this experiment is performed, the final product is actually lead. 

You can see that this experiment seems to prove the correctness 
of the phlogiston theory. Why? What mistake did the early sci- 
entists make in attempting to prove the phlogiston theory? 

For more than two centuries the phlogiston theory was considered 
to be an accurate explanation of burning, and many of the famous 
pioneers of modern chemistry, among them Priestley (prest'll) and 
Schcele (sha'l<?) were its ardent supporters. 

The first clue to the true explanation of burning. Kven though 
<the phlogiston theory seemed to be upheld by experiments similar 
to the one with lead, other experiments showed it to be false. In 
these experiments the fact that many substances increase in weight 
when burned could not be explained on the basis of the phlogiston 
theory, for according to it, all substances lose phlogiston when 
burned, and thereby lose weight. 



Lavoisier Priestley 

In 1774 Joseph Priestley, an English minister and amateur scien- 
tist, led the way toward a true explanation of burning by his dis- 
covery of a gas later named oxygen. Because of political and reli- 
gious persecution, he fled from England and spent his last years in 
Northumberland, Pennsylvania. Priestley showed that the gas he 
had discovered was present in air and was closely connected with 
burning. Let us trace the steps that led to the discovery of the true 
nature of burning. 

1) After his discovery of oxygen, Priestley, like a true scientist, 
did not keep his discovery to himself. While he was in Paris later 
on in 1774, he visited Lavoisier (la-vwa-zya') , the most eminent 
chemist in France, and told him about his discovery. Priestley's in- 
formation was a welcome addition to the many facts which Lavoi- 
sier had already collected regarding the nature of burning. 

2) Lavoisier examined carefully all the facts that he knew about 
burning. He pondered over them for months, trying to formulate 
an accurate theory that would explain burning and be in keeping 
with all the observed facts. 

3) Lavoisier was genius enough to use Priestley's work as the 
basis of a theory which would explain the age-old puzzle of burning. 
He suspected, and later advanced the theory, that when a substance 
burned, it increased in weight because it united with something that 
was present in air. Later this something was shown to be oxygen. 

4) Other scientists had noticed this increase in weight when sub- 
stances were burned in air. However, Lavoisier was the first to 
formulate a theory based on this fact. He also undertook a series 
of careful experiments to see whether their results would prove his 
theory to be correct. 


Lavoisier's classic 12-day experiment which explained burning. 

"I introduced four ounces of pure mercury into a [sealed] glass ves- 
sel," he wrote. "I lighted a fire in the furnace which I kept up con- 
tinually for twelve days. On the second day, small red particles 
already had begun to appear on the surface of the mercury." When 
most of the mercury had been converted into a red powder, he re- 
moved the glass vessel and its contents (which he had weighed before 
the experiment) and weighed them again. There was no increase 
in weight. 

Since the glass vessel was sealed, nothing had entered or escaped 
from it during the heating. Yet when he broke the seal he noticed 
that air rushed into the vessel. To him this inrush of air indicated 
that part of the air in the vessel had been used up during the heat- 
ing, and had left space for more air to enter. After air had entered 
the vessel, he weighed it once more and determined the increase in 
weight. He concluded that this increase in weight equaled the weight 
of something in the air in the vessel that must have combined with 
the mercury, forming the red powder. 

Lavoisier's inquiring spirit was not satisfied. He was a scientist in 
the most modern sense. He refused to jump to a hasty conclusion on 
the basis of a single experiment. He withheld drawing a conclusion 
until he had performed many more experiments. As a further pre- 
caution, he reversed his original experiment. He took the red pow- 
der of mercury and heated it to a higher temperature. He found 
that all of the red powder was changed back into mercury and that 
a gas was given off, which he found by a series of tests to be identi- 
cal with the oxygen gas that Priestley had discovered. Hence, he con- 
cluded that it was the oxygen in the air that was responsible for 
burning. Of all the substances he tried, he found none that could 
burn without oxygen. 

5) Burning, said Lavoisier, is the chemical union of a burnable 
substance with oxygen. Simple enough. No mysterious phlogiston, 
and the testimony of the most sensitive balance in Europe to sup- 
port his reasoning. Thus, Lavoisier discovered the true explanation 
of burning. 

6) Lavoisier repeated the original experiment using other sub- 
stances, including tin and sulfur. He found that the results of these 
experiments were fully in accord with his theory. In this way his 
theory was given further support. 


"Nothing happens without a cause/ 9 said Leucippus 2500 years 
ago. Ever since man first appeared on earth, he has been working 
constantly to find out the why of many natural occurrences. 

What are some of these occurrences that man has tried to explain? 
Although these are only a few, we might include: Why does a stone 
fall? What is fire? Why do some substances burn while other sub- 
stances do not? What makes thunder? Why can birds fly? 

In asking ourselves these questions and in answering them, we use 
several words that probably were not used by primitive man. These 
words, which appear later on in this paragraph, are words whose 
meanings have come into rather common use comparatively recently, 
as scientists measure time, perhaps within the last 25,000 or 50,000 
years. Let us examine the first question. Probably we would ask: 
What causes a stone to fall? By this we mean: What force causes a 
stone to fall, for we know that a stone will not fall unless some lorce 
acts on it, producing, we might say, an effect. We would answer the 
question: A stone falls because the earth pulls the stone toward its 
center (the force of gravity). As Leucippus (lu-slp'iis) implied, 
modern scientists and most modern people believe that every cause 
has an effect, and every effect has a cause. Such a relationship is 
known as a cause-and-effect relationship. 

Establishing a cause-and-effect relationship is not as simple as it 
might seem. Early man did not have the many tools and instruments 
which today we use so casually in finding cause-and-effect relation- 
ships. Consequently, in attempting to explain the causes of a certain 
effect, early man relied on what we would now consider magic, mys- 
ticism, and superstition, but later on man learned to establish these 
relationships by other methods. Then, too, it is sometimes difficult 
to establish a cause, because often several causes taken together pro- 
duce a single simple effect. Today many cause-and-effect relation- 
ships are clearly understood; but on the other hand, the causes of 
certain effects are not yet known, or even when known, are not thor- 
oughly understood. 

The method of deduction compared with induction. In explain- 
ing a natural occurrence, Aristotle (ar'Is-tot"l) , a well-known teacher 
and philosopher of ancient Greece, often made a bold and sweeping 
general statement. From this general statement he drew inferences 
and conclusions, which he thought applied in other similar cases. 
Aristotle's method is commonly known as the method of deduction. 

Francis Bacon, who lived almost 20 centuries later, was the first 
man to make popular another method of reasoning. By a process con- 
sisting of observation, collection of facts concerning the problem, 


formulation of a theory taking into account and explaining the 
observed facts, and verification of the theory by actual experiment, 
he formulated broad principles, sometimes called laws. Bacon's 
method is known as the method of induction, and today it is used 
widely. It is the pattern on which scientific method is based. To a 
great extent, it is responsible for the success of scientific method. 

While the method of induction is used very widely, deduction has 
played its part too in the development of science. The establishment 
of cause-and-effect relationships, usually by the method of induction, 
is perhaps the greatest function of science. 

What is scientific method? The method that Lavoisier used in 
reaching the first correct explanation of burning is an example of 
a pattern of action and thought used by scientists in their work. This 
pattern is known as scientific method. 

Scientific method may vary according to the nature of the problem 
to be solved and the tools available for solving it. In general, how- 
ever, the steps of the scientific method are represented by the six 
steps you have just traced. In brief, they may be stated as follows: 

The collection of all available facts related to a problem 

The open minded checking and examination of these facts 

The formulation of a working theory based upon these facts 
The testing of this working theory by experiments 

The formulation of a law, or principle, from the tested theory 
The use of this law in concrete and specific situations 

tr ' '. 

As you see, a scientific law, or principle, is a descriptive and ex- 
planatory statement, or generalization, that expresses what men have 
found to be accurate with respect to natural occurrences. 

Triumph of scientific method. Lavoisier's explanation of burning 
was not at once accepted. Indeed, for some time it met with bitter 
opposition. Those who believed in the phlogiston theory attempted 
to adjust their theory to fit the newly discovered facts, but this could 
not be done. Even so, Lavoisier realized how hard it would be to 


convince everyone of the truth of his own theory. He wrote: "I do 
not expect that my ideas will be accepted at once; the human mind 
inclines to one way of thinking, and those who have looked at na- 
ture from a certain point of view during a part of their lives adopt 
new ideas only with difficulty; it is for time, therefore, to confirm 
or reject the opinions that I have advanced. Meanwhile, I see that 
young men who are beginning to study the science (chemistry) 
without prejudice or preconceived notions no longer believe in 

New ideas, discoveries, and new theories as a rule must overcome 
tradition and prejudice, but this should not discourage those who 
introduce them. Tradition and prejudice are found not only in the 
field of the natural sciences physics, chemistry, astronomy, biol- 
ogy, and others but also in the social sciences government, eco- 
nomics, sociology, history, and others in which they are likely to 
be even more pronounced. 

Lavoisier's theory of burning finally triumphed, however. The 
accumulated evidence in its favor finally became so overpowering 
that scientists could believe nothing else. The experimental method 
of science had won over the strictly logical and theoretical method of 
the ancients. Twenty years later Lavoisier, who was an aristocrat, was 
beheaded during the frenzy of the French Revolution. "Until it is 
realized that the gravest crime of the French Revolution was not the 
execution of the king, but of Lavoisier, there is no right measure of 
values, for Lavoisier was one of the greatest three or four men 
France had produced." This statement, made by an eminent French- 
man, expresses the judgment of thinking men the world over. 

There are no authorities or dictators in science. There are only 
those persons who know what men up to now have discovered; that 
is, there are experts. In all of science there is no one who can say 
"This theory is true." But there are many men who can say "On the 
basis of what we now know, this theory seems to be sound." Scien- 
tific method is a democratic method. In most cases, its theories are 
the outcome of a pooling of facts discovered by many workers, a 
cooperative effort. 

Careful weighing and the law of the conservation of matter. To 
Lavoisier a balance was absolutely necessary. By carefully weighing 
all the substances entering into his experiments both before and 
after each experiment, he found that there was no loss of weight 
during the burning. The mercury plus the oxygen from the air 
weighed exactly the same as the red powder of mercury (mercuric 
oxide) . "One may take it for granted," he wrote, "that in every 



change there is an equal quantity of matter before and alter the op- 
eration." In chemical changes we can change the form, the state, or 
the composition of matter, but we cannot destroy matter itself. In 
physical changes we can change the ionn or the state of matter, but 
we cannot destroy the matter itself. 

"But/* you might answer, "a candle burns until i( is all gone. It 
becomes smaller and smaller and certainly weighs less and less." 
And you are right. But if we take the trouble to collect and weigh 
all the gases formed during the burning of the candle, we lind that 
they weigh more than the original candle. This increase in weight 
is due to the oxygen with which the burning candle has combined. 

"What of an oak tree," you might say. "It grows from a little 
acorn. Isn't matter created here?" It might seem so, but the fact that 
many things increase gradually in size and weight does not mean 
that matter has been or is being created. The oak tree does not come 
from seed alone. The cells from which the tree is made are built up 
chiefly from food materials taken out of the air, water, and soil. 
Parts of the air, water, and soil have been chemically changed and 
combined, forming living substances. 

Matter can be neither created nor destroyed. This fundamental 
law both of chemistry and of all science, is called the law of the con- 
servation of matter. 

1. Th a* o 

III* At 

th ttt 

th* the 

tilti on ttt 





Energy, too, can be neither created nor destroyed. So far we have 
considered mainly the changes in matter that take place in burning. 
But there are other changes that are equally important. When a 
substance burns, heat is liberated. This heat may be used to convert 
water in a boiler into steam. The pressure of this steam may then be 
used to turn a wheel, thus producing rotary motion. By connecting 
this wheel to a dynamo, electricity may be generated. This electricity 
may be changed into heat or light or magnetism, depending on 
whether it is sent through a toaster, a light bulb, or the electromag- 
net of, let us say, a buzzer or telegraph sounder. Or this electricity 
may be used to charge a storage battery. In charging the battery, the 
electricity produces a chemical change, which, on being reversed, 
yields electricity again. 

Evidently all these heat, electricity, the power to produce mo 
tion, light, magnetism, and the power to produce a chemical change 
can be transformed one into the other. All are capable of doing 
work, and all are forms of energy. In all energy changes, just as in 
all changes in matter, there is no loss, only transformation. Energy 
can be neither created nor destroyed. This is a fundamental law both 
of chemistry and of all science. It is called the law of the conserva- 
tion of energy. 

Two laws or one? Researches on the structure of matter and the 
nature of energy resulted in the atomic bomb. These researches, 

one of 

be of energy, follow 

in the 

discussed in Chapter 12, lead definitely to the conclusion that mat- 
ter and energy are but different forms of the same thing, and that 
matter can be converted into energy and energy into matter. 

As a result, the law of the conservation of matter and the law of 
the conservation of energy are no longer considered separate and 
distinct laws. Instead, they may be considered as different phases of 
a single law. Such a law would state that matter and energy can he 
neither created nor destroyed, but that each can be transformed 
into the other. In the transformation of matter into energy, matter 
disappears and becomes energy. In the transformation of energy into 
matter, energy disappears and becomes matter. 

Can any form of energy produce a chemical change? When pa- 
per is heated to its kindling temperature (see page 28) , it burns. 
Heat, one form of energy, produces a chemical change. When light, 
another form of energy, strikes a photographic film, it causes a chemi- 
cal change in the substances that coat the film. Thus, light also can 
cause a chemical change. When an electric current is passed through 
water, it splits the water into two gases, neither of which resembles 
water vapor. Thus electricity, too, is very effective in producing a 
chemical change. (The energy that was used to split the water ap- 
pears again as heat energy when the two gases are recombined, form- 
ing water.) From these experiments and from others, we know that 
many forms of energy bring about chemical changes. 

" '*' "*'- """~ 







The literature of chemistry is filled with romance. In it you 
Will find stories of human struggle and achievement, stories 
whose truth makes them all the more worthy to be read and 
remembered. Even the most interesting and significant parts 
of this literature cannot all be listed in this book. Neverthe- 
less, they are available, and it will be well worth your while 
to read as many of them as you can. The following is a selected 
list of books that deal with some of the topics in this chapter. 

Becker, Carl; Painter, Sidney; and Han, Yu-Shan. The Past 
that Lives Today, pp. 22-86. Silver Burdett Co., New York, 
1952. A fascinating story of early man, ancient civilizations 
and the science of the ancients is told here. 

Fabre, Jean H. The Wonder Book of Chemistry, pp. 6-69. 
Albert & Charles Boni, New York, 1922. A delightful account 
of elements, mixtures, and compounds. 

Jaffe, Bernard. Crucibles: The.Story of Chemistry, pp. 34-50. 
Simon and Schuster, New York, 1948. A simple account of the 
phlogiston theory. 

Somerville, John. The Way of Science: Its Growth and 
Method, pp. 93-113. Henry Schmnan, New York, 1953. A very 
simple illustration ot the steps ot scientific method. 


1. The three states of matter are solid, liquid, and gaseous. 

2. In a physical change, the original substance does not 
change into one or more other substances. 

3. In a chemical change, the original substances disappear 
and new substances are formed. 

4. An element is a substance that cannot be broken down, 
or decomposed, into a simpler substance by the ordinary types 
of chemical change. 

5. A chemical compound is a substance composed of two 
or more elements so combined that the elements can no longer 
be identified by their original individual properties. 

6. A mixture is a substance composed of two or more ele- 
ments or compounds that still retain their individual prop- 
erties after they have been thoroughly mixed. 

7. Burning is the chemical change in which a burnable 
substance unites with oxygen. 

8. The law of the conservation of matter states that mat- 
ter can be neither created nor destroyed. But matter may be 
changed from one form to another. 


9. The law of the conservation of energy states that energy 
can be neither created nor destroyed. But energy may be 
changed from one form to another. 

10. Recent researches prove that matter can be transformed 
into energy, and that energy can be transformed into matter. 

11. Establishing accurate cause-and-effect relationships is 
perhaps the greatest function of science. 

12. The steps in the scientific method include (1) the col- 
lection of all available facts related to a problem, (2) the 
open-minded examination of these facts, (3) the formulation 
of a working theory based upon these facts, (4) the testing of 
this working theory by experiments, (5) the formulation of a 
law, or principle, from the tested theory, and (6) the use of 
the law in specific situations. 

13. Blind acceptance of so-called "authorities," prejudices, 
and personal likes and dislikes have no place in the general 
pattern of action and thought of the true scientist. 


Group A 

1. (a) Which of the following are physical changes and 
which are chemical changes? (b) Give reasons for your 
answers. Souring of milk, molding of clay, digestion of food, 
drying of clothes, dissolving sugar in water, decay of fruit, 
freezing of mercury, photosynthesis (manufacture of starch in 
the leaves of plants from carbon dioxide gas and water) , 
erosion (the breaking up of rocks and soil by the action of air 
and water) . 

2. (a) What were the "four elements" of the ancients? 
(b) In what respects did their observations seem to support 
this theory? (c) What was one weakness of the way in which 
this theory was formulated? 

3. What are your reasons for thinking that water is a 

4. Classify the following as elements, compounds, or mix- 
tures: table salt, mercury, aluminum, paper, carbon dioxide, 
gold, silver, iron rust, sugar, sulfur, milk, brass, silver coin. 

5. Name a compound that contains hydrogen. 
4. Explain one method of telling a mixture from a com- 


7. Give briefly the main ideas of the "phlogiston theory." 

8. What was Priestley's part in discovering the true nature 
of burning? 

9. (a) Complete the following statement: When red mer- 
curic oxide is heated, it is changed into and (b) Is 

the change a physical or a chemical change? (c) Why? 

I . T . 

' i 

10. Describe Lavoisier's 12-day experiment on burning. 

11. Is Lavoisier's explanation the modern explanation of 

12. What part did the balance play in the development of 

13. State the six general steps in the scientific method. 

14. Distinguish between the method of science and the 
method ol the ancient teachers and philosophers. 

I ... 

15. (a) What is the law of the conservation of matter? 
(b) Is the disappearance of camphor balls in clothes an excep- 
tion to this law? (c) Explain. 

16. (a) What is energy? (b) Name three different forms of 

17. Give an illustration to show that each of the forms of 
eneigy named in your answer to question 16 may produce a 
chemical change. 

18. Assume that coal is the source of energy that lights your 
home (electric light). Make a list of the transformations of 
energy that occur, beginning with the burning of coal and end- 
ing with the lighted bulb. 

Group B 

19. Not many of the chemical elements w r ere discovered by 
Americans, (a) Which elements are these? (b) Can you sug- 
gest a reason why Americans have discovered so few? 

20. Aluminum is the most abundant metal in the earth. 
Tell why only in recent years aluminum has come into com- 
mon use. 

21. (a) Can one use scientific methods in fields other than 
science? (b) Explain your answer. 

22. (a) State four evidences of chemical action, and (b) give 
one example of each. 



1. Discuss with your teacher of history the problem of 
prejudice and tradition as obstacles to progress. Write a report 
on this topic using an illustration from American history or 
from the history of science. 

2. Find an article on some subject such as mental telepathy, 
communications from the spirit world, extrasensory percep- 
tion, astrology. Read the article carefully and write your own 
reaction as to whether you, using the methods of science, 
would accept the author's conclusions as scientific. 

3. Obtain a new photoflash bulb and weigh it on a sensitive 
balance. Weigh it again after it has been ignited. Do your 
findings uphold the law of the conservation ol matter? 

4. Show how you would use the method of science in solving 
some particular everyday problem which you have had to 
solve or which you will soon have to solve. 

5. With the aid of a medicine dropper, allow a drop of ink 
to fall into a tall glass of cold water. Observe what happens 
after approximately 5 minutes, 8 hours, 2-1 hours, and 2 days. 
Report to class with reference to physical and chemical 
changes, and any other conclusion you have drawn. 

6. Prepare a brief report on an expert in some field of sci- 
ence. Find out how he became an expert. How does an expert 
differ from an authority? 

7. On the first page of this chapter, there is a statement 
made by Lavoisier. Read it carefully. Report to your class on 
the importance of this statement. What are its implications 
for everyday living? 


. . . / procured a mouse and put it 
into a glass vessel containing the air 
(oxygen) from the red powder of 
mercury. Had it been common air, a 
full-grown mouse, as this was, would 
have lived in it about a quarter of an 
hour. In this air, however, my mouse 
lived a full half hour. 

Joseph Priestley, 1775 

A Sunday experiment by an English minister. On Sunday, the 
first of August, 1774, Priestley was working in his laboratory. He 
placed a red powder (mercuric oxide) in a bell jar so arranged that 
any gas which might be formed would pass out of the bell jar 
through a tube and be collected in a bottle. 

Instead of heating the powder over a flame, he used a large burn- 
ing lens to concentrate the rays of the sun on the powder. "I pres- 
ently found/' lie reported, "that air was expelled from it readily." 
But this result was not startling, because others before him had ob- 
tained gases by heating solids. 

A candle was burning in the laboratory. Wondering what effect 
this gas would have on a flame, he placed the candle in a bottle of 
it. Priestley reported in somewhat flowery words: "A candle burned 
in this air with an amazing strength of flame; and a bit of red-hot 
wood crackled and burned with a prodigious rapidity exhibiting an 
appearance something like that of iron glowing with a white heat 
and throwing out sparks in all directions.'* 

Only natural curiosity or perhaps chance led Priestley to experi- 
ment with the gas (or as he called it, air) . As Priestley, himself, said 
later, he had no idea what the outcome might be. 


Priestley was unable to explain what had happened. He was such 
a firm believer in the phlogiston theory that he did not associate this 
new gas with burning. As we learned in Chapter 1, it was Lavoisier 
who showed that Priestley's air (later called oxygen) is the element 
necessary for burning, thus solving a mystery that had baffled scien- 
tists for centuries. 

Half the earth is oxygen! Priestley's discovery of oxygen was a 
turning point in the development of chemistry. It is one of the 
strangest facts in history that this element, which surrounds us every- 
where and without which life is impossible, was not obtained pure 
until about 180 years ago. This fact is even more surprising when 
we realize that this one element, oxygen, is present on earth in quan- 
tities equal to the weight of all the other chemical elements put 
together. Sand and half of the different kinds of rocks on the earth 
are compounds of oxygen. Water contains almost 99 percent oxygen 
by weight, and air contains about 21 percent oxygen by volume. 

How oxygen is prepared in the laboratory. Because mercuric 
oxide, from which Priestley prepared oxygen, costs about one dol- 
lar a pound, it is too expensive to use in the laboratory. Instead, we 
obtain oxygen from a white crystalline solid called potassium chlo- 
rate. This compound is composed of three elements potassium, 
chlorine, and oxygen. By applying heat, the oxygen can be torn 
away and liberated as free oxygen. 

How to express the change by which oxygen is prepared. As po- 
tassium chlorate is heated, it yields potassium chloride plus oxygen. 
This is a chemical reaction which produces the chemical change rep- 
resented by the equation: 

Potassium (K) \ -\ 

^, , . /X; potassium 

Chlorine (Cl) \ -*> r ul . > -f oxygen 

,lv I chlorine J 75 

Oxygen (O) J J 

2KC1O 3 - 2KC1 + 30 2 

potassium chlorate > potassium chloride -f oxygen 

The chemical shorthand used to express the change will be fully 
explained later. The forms of energy that take part in the change are, 
as a rule, omitted from the equation. A chemical change is the re- 
sult of a chemical reaction. 

How the speed of this reaction can be increased. The method of 
preparing oxygen just described has one serious drawback. Unless 
a very high temperature is reached, oxygen is liberated very slowly. 
Someone discovered that if a small amount of powdered manganese 
dioxide, a black solid, is added to the potassium chlorate before 


heating, oxygen is liberated more quickly and at a lower tempera- 

At the end of the chemical reaction, that is, when oxygen is no 
longer given off, the same amount of manganese dioxide with which 
the experiment started remains. The weight of manganese dioxide 
has not been changed in any way. None of its oxygen has been lib- 
erated. Since manganese dioxide remains unchanged at the end of 
the reaction, it is not included in the equation. 

A catalyst changes the speed of a chemical reaction. Chemists 
have discovered that many chemical reactions can be speeded up or 
slowed down by placing a small quantity of certain substances in 
contact with the reacting materials. A substance that changes the 
speed of a chemical reaction is called a catalyst, or a catalytic agent. 
The catalyst itself may undergo some temporary change, but at the 
end of the reaction is present in the same state and quantity as at 
the beginning. Manganese dioxide, as used in the laboratory prepa- 
ration of oxygen, is a catalyst. However, manganese dioxide does 
not always act as a catalytic agent. It (see page 137) actually enters 
into certain reactions in which its composition is changed perma- 

We know little about the reasons that catalysts act as they do. 
We do know, however, that many vitally important chemical re- 
actions take place too slowly, and hence uneconomically, except in 
the presence of certain catalysts. Research on the nature of catalysis 
is now being carried on in laboratories throughout the world, and we 
will know much more about it before very long. 

The presence of a catalyst in a reaction is sometimes indicated by 
writing the catalyst over the arrow in the equation. In the reaction 
just discussed, the presence of the catalyst, manganese dioxide, might 
be indicated thus: 

MnO 2 
2KC1O 3 > 2KC1 + 3O 2 

How oxygen gas is collected. In preparing oxygen, a mixture of 
potassium chlorate and manganese dioxide is put in a test tube and 
heated over a bunsen burner (see illustration page 27) . Connected 
to the test tube is a delivery tube which reaches into a bottle that 
has been filled with water and placed mouth downward in a pan of 
water. As the potassium chlorate is heated, it is broken down, or 
decomposed, forming potassium chloride and oxygen. 

The potassium chloride is a solid and remains in the test tube. 
The oxygen is a gas and passes through the delivery tube into the 

2. of if 

ii the 

water-filled collecting bottle from which it displaces water. This 
method of collecting gases is called the displacement of water method, 
and was used by Priestley. Priestley also collected gases by the dis- 
placement of mercury when such gases dissolved in water but not 
in mercury. If a gas is collected by the displacement of a liquid, 
chemists say that the gas is collected over the liquid. Thus, Priestley 
collected oxygen over water. 

Physical properties of oxygen. In discussing the physical properties 
of a gas, we usually consider five characteristics: (1) color, (2) odor, 

(3) weight compared with the weight of an equal volume of air, 

(4) the ease with which it may be changed into a liquid, and (5) its 
absorption by water. This absorption by water we call its solubility 
in water. Just as solids, such as sugar and salt, disappear as solids 
when stirred in water and are distributed uniformly throughout 
the liquid, so gases, such as air and oxygen, when passed through 
water, may likewise be absorbed by the water. In the case of some 
gases, taste is considered also. 

Oxygen is colorless and has no odor. It is slightly heavier than 
air. It is slightly soluble in water. Under normal conditions, that 
is, at a temperature of 18 degrees centigrade and a pressure of one 
atmosphere, or 760 millimeters of mercury,* about four quarts of 

* Scientists often use two terms to describe temperature and pressure, normal 
conditions and standard conditions. Standard conditions mean a temperature 
of C. and a pressure of 760 mm. of mercury, that is, one atmosphere (atm.) . 
Normal conditions mean, in effect, roorn temperature and the pressure of 1 at- 
mosphere. Throughout this book a pressure of 760 mm. is assumed when staling 
temperatures at which gases liquefy or liquids solidify. A millimeter (mm.) is 
a small unit of length in the Metric System. Numerically it is equal to 0.001 of 
a meter, a larger unit of length equal to 39.37 inches. For an explanation of 
metric units and temperature scales, see pages 654 and 644 respectively. Mcttic. 
units of measurement are used by scientists in all countries. 


oxygen gas will dissolve in 100 quarts of water. It is hard to change 
oxygen to a liquid. At a temperature of about 183 degrees below- 
zero centigrade ( 183C.) and under a pressure of 760 millimeters, 
oxygen is converted into a pale blue liquid, which can be attracted 
by a magnet. At 219C. it changes into a bluish-white solid. 

Chemically, oxygen is an active element. Because oxygen combines 
with almost all other elements, forming compounds called oxides, it 
is considered to be very active chemically. For example, when iron 
is exposed to oxygen, rust, which is an oxide of iron, is formed. 

Iron + oxygen > iron oxide 
4Fc + 3O 2 -> 2Fc 2 O 3 

If iron is first heated until it glows and then placed in a bottle of 
oxygen, the chemical reaction is so vigorous that the iron burns 
brilliantly, throwing off sparks of glowing iron oxide. This sur- 
prising spectacle actually is iron burning in oxygen, just as a piece 
of paper burns in air! 

Differences between slow and rapid oxidation. When a sub- 
stance element or compound combines with oxygen, new sub- 
stances are formed. This chemical union of a substance with oxygen 
is called oxidation. Rapid oxidation, such as the burning of coal, is 
accompanied by noticeable heat and light. During slow oxidation, 
such as the rusting of iron or the decaying of wood, no light is given 
off nor can we easily detect the heat because it is given off so slowly. 
Delicate measurements, however, have proved beyond doubt that 
the amount of heat energy liberated is the same whether the oxida- 
tion of a substance takes place slowly or rapidly. 

The soft cold light of the firefly, and the glow of some fungi and 
bacteria are caused by the oxidation of a complex chemical com- 
pound, luciferin, which they produce. 

Why some substances catch fire more easily than others. Some 
substances catch fire at low temperatures, but others require ex- 
tremely high temperatures in order to burn. Every substance must 
be raised to a certain definite temperature before it can combine 
with oxygen at such a rate that the heat produced is sufficient to keep 
the substance burning without the addition of more external heat. 

In starting a coal fire, we often begin by burning paper, which 
sets fire to kindling wood, which sets fire to the coal. The heat given 
off by the burning paper causes the wood to catch fire; the heat given 
off by the burning wood in turn causes the coal to catch fire. The 
lowest temperature at which a substance catches fire and continues 
to burn is called the kindling temperature of that substance. We 

Hitreau of Mintt, U.S. Depf. of Intu 

A coal-dust explosion issues from the U.S. Bureau of Mines* 
Experimental Coal Mine in Pennsylvania. Such test ex- 
plosions are part of the Bureau's safety research program. 

say that a substance which is easy to set on fire has a low kindling 
temperature, and a substance which is difficult to set on fire has a 
high kindling temperature. 

A substance may have different kindling temperatures, depending 
upon the size of the particles into which it is divided, that is, upon 
its state of subdivision. A solid piece of iron has a high kindling 
temperature, but powdered iron, because of the large surface that 
is exposed to the oxygen of the air, can be made to burn readily 
in air. Many dust explosions in flour mills, starch factories, grain 
elevators, and coal mines are caused by the very rapid oxidation of 
explosive mixtures of air and finely divided materials. A spark 
resulting from static electricity or friction often sets off the explosion. 

Smut dust and air form an explosive mixture which may be 
ignited by static electricity during threshing operations. Costly fires 
of this kind have been so widespread that the United States Depart- 
ment of Agriculture has issued a bulletin explaining how to prevent 


The explosion of a mixture of coal dust and air has been used 
in one type of internal-combustion engine. 

Substances such as asbestos, brick, concrete, and marble never 
catch fire because they are already completely oxidized. 

Spontaneous combustion. Several years ago there were widespread 
floods in the Ohio Valley. The lower parts of thousands of haystacks 
in the Valley were soaked with water. As the flood waters receded, 
farmers were pu//led when some of the haystacks began to catch fire. 

The explanation of this phenomenon is simple. During respiration 
of living plant cells, food materials slowly oxidize and heat is given 
off. This oxidation is speeded up by the presence of a small amount 
of moisture. The hay itself also slowly oxidi/es, liberating heat. As 
the temperature rises, the rate of oxidation also increases. The heat 
slowly accumulates, and when the kindling temperature of the hay is 
finally reached, it bursts into flame. Materials catching fire in this 
way are said to undergo spontaneous combustion. (Combustion 
refers to any chemical reaction which produces heat and light. 
Burning is only one kind of combustion.) 

Fires have also been caused by painters' rags saturated with linseed 
oil. As the linseed oil slowly oxidizes, heat is given off. Unless there 
is a free circulation of air to carry away this heat, the oily rags may 
become hot enough to catch fire. Thus, you see why oily rags should 
not be kept in poorly ventilated closets. Finely divided coal in the 
closed hold of a ship, or in a poorly ventilated boiler room must be 
sprinkled with plenty of water from time to time to prevent the 
accumulation of heat from slow oxidation. 

Phosphorus, an element that burns spontaneously. Ancient alche- 
mists spent most of their time looking for the philosopher s stone, 
which they believed would change lead into gold. In 1669, while 
searching for the philosopher's stone, Hennig Brand, an alchemist 
of Hamburg, obtained a new and strange chemical substance from 
urine. It had so low a kindling temperature that on exposure to 
air it caught fire immediately and burned, forming a white powder 
(phosphorus oxide) . Brand made a famous tour of Europe, exhibit- 
ing this unusual substance. Today we know it as white phosphorus, 
a soft, waxy element, now obtained by a chemical process from bone 
deposits. At one time it was used in the manufacture of matches 
(see page 296) . This element is sometimes referred to as yellow 

To keep white phosphorus from catching fire, it must be stored 
under water. Do not touch it with your bare fingers, for white 
phosphorus will cause severe burns that heal very slowly. 



Upon exposure to air, white 
phosphorus ignites sponta- 
neously in a violent reaction. 
Note the protective clothing 
worn by the demonstrator. 

Monsanto Chemical 

You could not live without oxygen! One of the remarkable ex- 
periments which i'riestley performed showed that a mouse, placed 
in a bell jar, lived twice as long in pure oxygen as in "common air." 
This, Priestley was unable to explain; but today we know that the 
chief chemical change that goes on in the body of any living animal 
is slow oxidation. 

Man obtains his supply of oxygen by breathing. When he inhales, 
oxygen is taken into his lungs. This oxygen passes through the walls 
of the lungs and is absorbed by the red cells of the blood, which 
carry it to all parts of the body. In every living cell or body dim, 
oxidation takes place, liberating heat and other forms of energy. This 
slow, steady oxidation is like a tiny flame which keeps life going. 
Without oxygen, the flame is snuffed out and life is extinguished. 

With the exception of a few very low forms of life, all living 
things take oxygen from the air. This oxygen is in the /rvv, or uiic.oin- 
bined, state. That is, the oxygen is not chemically united with any 
of the other substances in air. Fish obtain their supply ol oxygen, 
from the air that is dissolved in water. 


Adapted from drawing by Linde Air Products Company 

Fig. 3. Oxyacetylene torch. What is the 
function of the expansion chamber? 

Industry uses great quantities of oxygen. Commercial production 
of pure oxygen in the United States is more than 25 billion cubic 
feet a year. Of the oxygen produced, it is estimated that more than 
95 percent is used in cutting and processing steel and in welding 
metals, such as aluminum and steel, by means of oxyacetylene and 
oxy hydrogen torches. 

It was Priestley who first thought of using oxygen to produce high 
temperatures. He found that blowing pure oxygen on a piece of 
glowing wood would cause it to burn furiously. A few years later 
an American scientist, Robert Hare, of Philadelphia, put this dis- 
covery of his friend Priestley to practical use by inventing the oxyhy- 
drogen torch, or blowpipe. 

The oxyhydrogen torch consists of two tubes, one inside the other. 
Hydrogen gas passes through the outer tube and is ignited at the tip 
of the torch. Pure oxygen passes through the inner tube and the 
mixture of the two gases burns at the tip of the torch with an ex- 
tremely hot flame, about 2400C., a temperature much higher than 
the melting point of iron. 

The oxyhydrogen torch was never widely used. Instead, the oxy- 
acetylene torch is used. With the oxyacetylene torch, a flame tem- 
perature of over 3300C. may be easily produced. The oxyacetylene 
torch is similar in principle to the oxyhydrogen torch, but acetylene 

Fig. 4. Blast lamp. Similar in principle 
to the oxyhydrogen torch except 
that compressed air and any fuel 
gas are used instead of oxygen and 
hydrogen. The lamp is used by glass 
blowers and jewelers. 






gas is used instead of hydrogen (see Fig. 3) . The oxygen used in 
the torch is stored under high pressure in strong steel cylinders. 
The acetylene, however, is not under high pressure, but is dissolved 
in a liquid called acetone. In almost every automobile service garage, 
oxyacetylene torches, with their accompanying cylinders ol oxygen 
and acetylene, may be seen ready for use. 

The oxyacetylene torch is an important industrial tool. It is used 
to weld, cut, and clean metal. It is also used in heat treating sur- 
faces of metal machine parts to make them more wear resistant. 

Oxygen rusts and derusts steel. In the presence of air and minute 
quantities of water vapor, steel rusts. In rusting, oxygen from the 
air unites slowly with the metal, forming a brown, scaly oxide of 
iron. At the high temperatures used in the making of steel, rust lorms 
very rapidly and is a serious problem. As red-hot steel is carried to 
the rolling mills, it becomes covered with seams of iron oxide. 

These surface imperfections are removed by a process called 
torch-deseaming, or scarfing. In the scarfing operation, the llame of 
the oxyacetylene torch is directed onto the hot steel. The surface 
of the steel is quickly oxidized to a depth of about one quarter of an 
inch. The iron oxide falls off readily, leaving an unblemished sur- 
face. Scarfing may be done by hand or by special machines. 

(left) Red-hot steel slab passes through scarfing machine 
in which oxyacetylene torches remove surface defects, 
(right) Welding metal plates with an oxyacetylene torch. 

Linde Air Products Company 



Because of the high tem- 
perature of the torch, an 
oxyacetylene weld is 
smooth and strong. 

/.nidc Air Product* Company 

Oxygen saves lives. Priestley also discovered another use for oxy- 
gen. Alter inhaling oxygen, he wrote: "My breath felt peculiarly 
light and easy. It (oxygen) may be peculiarly salutary to the lungs in 
certain cases where the common air is not sufficient." 

Today, pure oxygen is administered to persons with pneumonia 
and in other cases where the respiratory system cannot function at 
its normal rate. Usually about two gallons of oxygen are administered 
per minute. Since air contains only about 20 percent pure oxygen, 
a patient who is weak, or whose lungs are congested or partially 
destroyed, can satisfy the oxygen requirements of his body by breath- 
ing a much smaller volume of air that is rich in oxygen than he 
would normally require of ordinary air. Oxygen is usually admin- 
istered by means of an oxygen tent, a canopy which fits over the 
patient's bed. Pure oxygen is introduced at such a rate that the air 
inside the canopy always contains from 45 to GO percent oxygen. 

Oxygen in determining basal metabolism. Pure oxygen is used by 
physicians in determining the rate at which a person's food supply 
is oxidized while the person is at rest. This is known as his rate of 
basal metabolism. This rate is obtained by measuring the volume of 
oxygen consumed by the person at rest during a short interval, 
usually eight minutes. 

From these data the number of liters* of oxygen consumed per 
minute may be calculated easily. This number is then compared 
with the basal metabolic rate for a normal person of the same age, 
sex, height, and weight. Persons in normal health use oxygen at a 

* A liter (I.) is a unit of capacity (volume) in the Metric System. It is 
slightly larger than a U.S. liquid quart. 



standard rate. In certain diseases, the patient's rate is higher and 
in others it is lower than known standards for persons in good 
health. For example, a high basal metabolic rate always accompanies 
an overactive thyroid (see pages 147, 581). A low basal metabolic 
rate may be an indication of an underactive thyroid. 

Oxygen flies high. Aviators and mountain climbers who ascend to 
high altitudes where the atmosphere is very thin must carry supplies 
of oxygen. Otherwise their senses become dulled and they are likely 
to lose consciousness. The United States Air Force requires the use 
of oxygen at altitudes above 10,000 feet in the daytime and from the 
ground up at nighttime. Airliners flying at altitudes ranging from 
15,000 to more than .SO, 000 feet contain equipment to keep the oxy- 
gen concentration inside their cabins only slightly less than that at sea 
level. The cabins are airtight and are pressurized by means of pumps, 
so that the inside pressure does not become uncomfortably low. 

Rescue parties entering mines and buildings in which dangerous 
gases are present carry oxygen-breathing apparatus. 

Air Photographic and Charting Service, U.S. Air Force 

High in the strato- 
sphere, this pilot is de- 
pendent upon a con- 
tinuous supply of pure 


Some other uses of oxygen. Oxygen is used in the photoflash lamps 
employed in photography. These lamps look like ordinary electric- 
light bulbs, but they are filled with oxygen and aluminum foil. When 
an electric current is passed through the filament, the aluminum foil 
is raised to its kindling temperature. It ignites and oxidizes with 
a blinding flash. This Hash lasts only about ^ of a second. 

Pure oxygen has also been introduced in the manufacture of steel, 
synthetic gasoline, and fuel from underground coal (see pages 335, 
3(50, 417). 

How is such a huge amount of oxygen prepared? It is not sur- 
prising to find that the two main sources of oxygen are air, the most 
abundant mixture containing oxygen, and water, the most abundant 
compound ol oxygen. Nearly all commercial oxygen is obtained 
from liquid air by first lowering the temperature of the air until it is 
changed to a liquid. The preparation of oxygen from liquid air is 
discussed on page 100. Less than one percent of the oxygen produced 
commercially is obtained by decomposing water by means ol an 
electric current. This process, called the electrolysis of water, is de- 
scribed on page 62. 

How can we test for oxygen? When Priestley prepared oxygen 
from mercuric oxide, he tested the gas by placing burning and 
glowing substances in it. In each case, the substance burned more 
vigorously. This method is still used to identity oxygen. A glowing 
splint or splinter ol wood is thrust into a bottle oi the gas. Such a 
splint placed in a bottle of oxygen bursts into flame at once. 

No other odorless gas will cause a glowing splint to burst into 
flame in this way. Hence, we can distinguish oxygen from any other 
odorless gas by this simple procedure. We call such a method of 
identifying a substance a test for that substance. Chemists have 
devised hundreds of tests, which they use in identifying many other 
pure substances. 

Ozone, the active. Ten years after Priestley's discovery of oxygen, 
another gas which possessed a peculiar odor and which, unlike 
oxygen, tarnished mercury under normal conditions, was reported. 
But it was not until 1810 that Schoenbcin (shun'bfn) isolated this 
gas and called it ozone, from the Greek word meaning to smell. Its 
sharp odor is noticeable around electric machines in operation. 

O/one is a pale blue gas, one and one-half times as heavy as 
oxygen. It is even less soluble in water than oxygen but is more 
active chemically. It is a strong oxidizing agent. That is, it is a 
substance which readily supplies oxygen for chemical union with 
another substance. 

/. -' 

current L '" 11 ' 

Fig. 5. A continuous-process ozone tube. Dry air enters 
at lower left, passes through the brush discharge of high- 
voltage current. Part of the oxygen of the air is converted 
to ozone which leaves through pipe at right. 

Ozone (written O 8 ) is prepared by passing electric discharges 
through cither dry air or oxygen (written O.,) . About eight percent 
of the oxygen is converted into pure o/.one, although slightly larger 
yields are obtained if the temperature is kept low, or if the process 

is continuous. 

3 volumes of oxygen 
30 2 

- 2 volumes of ozone 
20 3 

Ozone is unstable; it changes back to oxygen quickly, two volumes 
of ozone changing into three volumes of oxygen. It cannot be stored 
and must be produced at its point of use. 

What is allotropy? As we have just learned, oxygen exists in two 
forms: ordinary oxygen and o/one. The existence in the same 
physical state (both oxygen and o/.one are gases) of two or more 
forms of the same element is a phenomenon called allotropy 
(0-lot'r6-pi) . The various allotropic forms of an element have differ- 
ent physical and chemical properties. 

The cause of allotropy is not yet completely understood. Hut we 
know that it is caused in part by differences in the arrangement of 
the atoms and in the amount of energy in the various allotropic 
forms. This can be seen easily by referring to the way in which o/one 
is produced from oxygen. When an electric current discharges 
through oxygen, the electric energy changes oxygen into ozone, 
which, as you would expect, possesses more energy than ordinary 
oxygen. O/one, on changing back into oxygen, liberates this energy 
in the form of heat. During the change from oxygen to ozone, or 
from o/one to oxygen, no energy is destroyed, nor is any energy cre- 
ated. The change from one allotropic form of oxygen to the other 
is an excellent example of the law of the conservation of energy. 



Allotropy is not confined, of course, to oxygen. For example, the 
element phosphorus occurs in two allotropic forms: white phos- 
phorus and red phosphorus. White phosphorus melts at 44C., is 
poisonous, soluble in carbon disulfide, and has a very low kindling 
temperature. Red phosphorus is nonpoisonous, insoluble in carbon 
disulfide, and has a higher kindling temperature than white phos- 
phorus. Red phosphorus is also heavier and less active chemically 
than white phosphorus. Under proper conditions, red phosphorus 
may be changed into white phosphorus, and vice versa. 

Ozone "burns up" germs. Because of its extreme chemical activity, 
o/one is used to a limited extent in purifying water. It kills bacteria 
and other microorganisms in water by oxidi/ing them, literally burn- 
ing them up. Other organic: materials present are also oxidized. 
About one gram* of ozone will purify a cubic meter of water. 

Ozone destroys odors. Because o/one is an excellent oxidi/ing 
agent, it is used also in purifying air in homes, refrigerators, tunnels, 
and zoos. Small ultraviolet lamps change the oxygen in air to ozone 
which clears away bad odors. Because of the increasing abundance of 
low-cost electricity, it seems possible that o/one may be used more 
widely in the removal ot unpleasant tastes and odors from water 
than it now is. 

Ozone helps screen the earth. Ozone is present in the layers of 
the atmosphere, about 30 miles above the surface of the earth. Sci- 
entists believe that this region which is rich in o/one acts as a screen 
that protects life on earth from the harmful effects of too much ultra- 
violet light from the sun. 

What part does chance play in scientific discoveries? Some years 
after his discovery of oxygen, Priestley commenting on this memo- 
rable occasion said: "I can not at this distance of time recall what it 
was that I had in view in making this experiment, but I had no 
expectation of the real issue of it. If I had not happened to have had 
a lighted candle before me, I should probably never have made the 
trial, and the whole train of my future experiments relating to this 
kind of air might have been prevented. More is owing to what we 
call chance than to any proper design or preconceived theory." 

Chance may have played some small part in leading Priestley to 
make his experiments. It seems likely, though, that he failed to take 
into account the consuming natural curiosity, always present in true 
scientists, which literally forced him to make his experiments. No 

* A gram (g.) is a small unit of weight or mass in the Metric System. One 
thousand grams are equal to a kilogram, which weighs slightly more than 2.2 
pounds. For a definition of mass, see page 157. 


doubt one of the tests he would eventually have made on any gas 
was to see if it would burn. If the lighted candle had not been 
present at that particular moment, he undoubtedly would have 
lighted one later on. Probably what Priestley meant by the statement 
was simply that he had no specific purpose in mind in making the 
experiment. But having no specific purpose in mind and attributing 
the results to chance are not the same. Nevertheless, such confusion 
exists among scientists even today. 

Other men were performing experiments similar to those made by 
Priestley (among them Scheele) , and we might say that the time was 
ripe for these discoveries. If Priestley had failed to make them, no 
doubt some other experimenter soon would have made them. 

Today carefully planned research has speeded up advances in 
science. Chance now plays a very slight role in the development of 
chemistry. Surely we cannot regard as "chance" the keenness of 
mind to appreciate the significance ol, and to follow up by intelligent 
experiment, a clue furnished by some unforeseen event. 

Progress in science and changes in society are closely interrelated. 
Most of us, although we live in a world ol science, have strange 
notions about what scientists arc like. We picture the research man 
as a lone, mysterious genius who locks himsell up in his laboratory 
away from the world, and attempts to solve some abstract scientific 
problem. We sec him emerging triumphant, after weeks or perhaps 
years, with some great discovery. Most of us have an idea that the 
efforts of the man of science arc influenced very little by the society 
in which he lives. Nothing could be farther from the truth. The 
scientist is influenced by society and society in turn is influenced 
by the scientist. 

The steam engine, for example, was developed out of the social 
needs of the eighteenth century. With the Industrial Revolution, 
which began in Birmingham, England, came a need for more iron to 
make the machinery so much in demand. In making iron, wood char- 
coal had been used, but England's forest reserves had been seriously 
depleted by the use of timber in the many ships swallowed up in 
naval wars and commercial ventures. As a result, coal had to be sub- 
stituted for charcoal, and coal mines, abandoned because they were 
Hooded, had to be reopened. This meant draining mines, and James 
Watt improved Newcomen's steam engine for this purpose. 

At about the same time, the problem of burning was reinvesti- 
gated, and Priestley's investigations (which led to his discovery of 
oxygen) were closely related to society's need for more information 
concerning the extraction of metals. 




Becker, Carl. Modern History, pp. 581-607. Silver Burdett 
Co., New York, 1952. "How Science Gave Men Machines to 
Work for Them, and How the Machines Changed the Con- 
ditions under Which Men Had to Live and Labor." Very 

Ficklen, Joseph B. "Dust Explosions." Journal of Chemical 
Education, 'March, 1942, pp. 131-134. Published by the Amer- 
ican Chemical Society, Easton, Pa. Editorial office: Metcalf 
Chemical Laboratory, Brown University, Providence, R.I. An 
interesting and well-illustrated article. 

Friend, J. N. Man and the Chemical Elements. Scribners, 
New York, 1953. The author rambles along into many inter- 
esting sidepaths and offers sidelights such as the origin of 
gibberish from the alchemist Geber. 

Partington, J. R. A Short History of Chemistry, pp. 110-120. 
The Macmillan Co., London, 1939. The fascinating story of 
Priestley's discovery of oxygen. 


1. A catalyst is a substance that changes the speed of a 
chemical reaction without being itself permanently changed. 

2. Oxidation is the chemical union of a substance with 
oxygen. In slow oxidation, neither light nor noticeable heat is 
liberated. In rapid oxidation, light and noticeable heat are 
evolved. Burning is rapid oxidation. 

3. The kindling temperature ot a substance is the lowest 
temperature at which that substance catches fire and con- 
tinues to burn. 

4. Spontaneous combustion is a burning started by the heat 
accumulated during slow oxidation. Combustion is any chem- 
ical action that liberates heat and light. 

5. Finely divided powders form explosive mixtures with air 
because of the extremely large surfaces exposed to oxygen. 

6. Elements may occur in two or more varieties, or allotropic 
forms, in the same physical state. These allotropic forms differ 
in their properties because of differences in the arrangement 
of their atoms and differences in the amount of energy pos- 
sessed by particles of each allotrope. 

7. Today carefully planned research has speeded up ad- 
vances in science, and chance discoveries play a smaller part 
in the development of chemistry than they once did. 

8. Scientific progress and social changes are interrelated. 



Group A 

1. Make as large a list as you can of substances containing 
oxygen combined with other elements. 

2. From what chemical compound was O 2 first prepared 
by Priestley? 

3. From what compounds is O 2 usually prepared in the 

4. (a) What is a catalyst? (b) Illustrate your answer. 

5. What assurance have we that the MnO., used in the lab- 
oratory preparation of O 2 acts as a catalytic agent? 

6. What happens to the KCl that is formed by the de- 
composition of KC1O 3 ? 

7. Why is it possible to collect O 2 by the displacement of 

8. Priestley for a time collected gases by the displacement 
of mercury. Why do we not use this method in collecting O.,? 

9. List five physical properties of O.,. 

10. Discuss the most important chemical property of O 2 . 

11. (a) What is slow oxidation? (b) rapid oxidation? (c) Il- 
lustrate each. 

12. What is the kindling temperature of a substance? 

13. Devise an experiment to show that air is necessary for 

14. Name four substances whose kindling temperatures are 
lower than that of coal. 

15. Why is air removed from an electric-light bulb? 

16. It is not as easy to burn a 500-page book whole as it is to 
burn the same book a page at a time. Explain. 

17. Why is asbestos used in making theater curtains? 

18. (a) What are the conditions necessary for spontaneous 
combustion? (b) Show how these conditions work out in the 
spontaneous combustion of (1) finely divided coal in the en- 
closed hold oi a ship, and (2) moist hay in a hayloft. 

19. Why may the presence of dust in the air of a Hour mill 
cause a frightful explosion? 

20. Make a list of the uses of O r 

21. Distinguish between combustion and burning. 


22. What is an oxidizing agent? 

23. Describe the greatest industrial use of O 2 . 

24. O 2 rusts steel. What part does O 2 play in derusting steel? 

25. How do oxygen tents aid patients suffering from pneu- 
monia and other respiratory diseases? 

26. (a) What is basal metabolic rate? (b) How is it de- 
termined? (c) What is its significance? 

27. How does O 2 play an important part in aviation? 

. . T . . . 

- I - 

28. State the two most important industrial .sources of O . 

29. What three factors must be considered in selecting the 
method for preparing large quantities of a substance? 

30. (a) Describe a test tor O 2 . (b) What does the word test 
mean to a chemist? 

31. O 2 and O 3 can be changed easily from one to the other. 
(a) How? (b) What great law does this illustrate? 

32. How is O 3 prepared? 

33. How does O 8 differ from O 2 in physical properties? 

34. By means ot an illustration, explain the meaning of 
a I lot ropy. 

35. What are the chief uses of O 3 ? 

36. Write word-equations for the burning of (a) coal (car- 
bon) , (b) iron, (c) phosphorus, and (d) sultur. 

Group B 

37. By using mercury Priestley was able to collect a number 
of gases that had escaped the attention of other scientists. 

38. Science today depends less upon chance discoveries than 
it has in the past. Explain. 

39. Window curtains behind a fish bowl filled with water 
caught fire. Was this spontaneous combustion? Explain. 

40. Why is O, passed through the inner tube ot the oxyhy- 
drogen torch rather than through the outer tube? 

41. Because ot oxidation, linseed oil hardens when exposed 
to air. How is this chemical reaction speeded up? 

42. Fishes die in air, man drowns in the sea. Explain. 

43. Blow against a burning candle and it goes out. Blow on 
the slowly dying embers of a fire and they burn more actively. 

44. Why is green, or moist, hay more susceptible to spon- 
taneous combustion than dry hay? 


45. Illustrate the statement "progress in science and changes 
in society are closely interrelated" by an example not men- 
tioned in this chapter. 

46. Coal dust is sometimes shoveled into a burning coal 
furnace. Why? 


1. Have you ever heard of pure oxygen being administered 
to athletes before a strenuous game? Consult the football or 
track coach in your school or in a nearby college. What do 
you conclude? Explain. 

2. Organize a small group in your chemistry class and with 
the help of your teacher set up apparatus such as Priestley first 
used when he obtained oxygen. A large burning glass can be 
borrowed from the physics department. Demonstrate the ex- 
periment before your class. 

3. Watch your mother put up fruit and jam in jars. De- 
scribe her procedure, and explain why the food is heated, 
and why a layer of paraffin is placed over the food before the 
jar is sealed. 

.4. Consult your doctor and a good book on first aid, and 
write a report on the diagnosis and treatment of first-, second-, 
and third-degree burns. 

5. Read the story of a recent scientific discovery, (a) From 
the account you read, what part do you conclude that chance 
plays in scientific discoveries? (b) Why may the scientist 
himself believe chance plays a more important part than it 
actually does? 

H Y D R O G E 


. . . Cavendish was almost passion- 
less. An intellectual head thinking, 
a pair of wonderful acute eyes ob- 
serving, and a pair of skilful hands 
* . experimenting or recording are all 

that I realize in reading his memoirs. 
Cavendish did not stand aloof from 
other men in proud spirit, he did 

> , , so conscious of his inferiority, not 

boasting of his excellence. 

Dr. George Wilson, 1851 

An eccentric man of science. While Priestley was performing his 
immortal experiments, another Englishman was puttering around 
in his palace laboratory on the wandering trail of phlogiston. This 
man was Henry Cavendish, one of the most eccentric persons in the 
whole history of science. It was said that he was "the richest among 
the learned and the most learned among the rich." Although the 
richest man in all England, he shut himself up in his private labora- 
tory and spent more than 60 years tracking down many of the secrets 
of nature. 

In addition to his discoveries in chemistry, Cavendish was inter- 
ested in physical problems. His work in the fields of heat and elec- 
tricity was of highest rank; later it was followed up by the work of 
other eminent scientists such as Joseph Black and Michael Faraday. 
The elusive trail of phlogiston. Two hundred and fifty years 
before the work of Cavendish, Paracelsus (par-d-sel'sus) of Switzer- 
land had noticed bubbles of gas rise from an acid into which iron is 
dropped. Although he found that this gas burns, he carried his 
investigations no further. Then came Cavendish to whom the search 
for truth was the ruling motive of life. He too noticed the gas 
evolved when zinc or iron is dropped into an acid and went to work 



to investigate this phenomenon. He collected the gas carefully, and 
made a thorough study of it. He named it inflammable air because it 
burned. He thought he had obtained phlogiston itself. 

Strangely enough, the discovery of this gas, coupled with the dis- 
covery of oxygen, paved the way for the complete overthrow of the 
phlogiston theory and the establishment ol l.axoisier's true explana- 
tion of burning. Though Priestley died still believing in the phlo- 
giston theory, Cavendish, when the discussion over Lavoisier's new 
chemistry became very heated, gave up his active interest in chemistry 
and turned to the problem of determining the weight of the earth. 
He said that he had no patience with squabbles, that he was inter- 
ested in experimentation, not controversy. 

How is hydrogen produced in the laboratory? Cavendish's lab- 
oratory method of preparing hydrogen is still used. Zinc is placed 
in a generator, as shown in Fig. 6, and dilute hydrochloric or sul- 
turic acid is poured over it through a thistle tube. Bubbles of hydro- 
gen gas form at once, and heat is generated. The hydrogen is col- 
lected in the same way as oxygen, that is, by the displacement of 
water. Why is hydrogen not collected by the displacement of air? 

The chemical change that takes place is represented as follows: 

Zinc -4- hydrochloric acid 
Zn + 2HC1 

> zinc chloride 4- hydrogen 
ZnCl 2 4- H 2 

Hydrochloric acid is a compound of hydrogen and chlorine. The 
zinc takes the place of, or replaces, the hydrogen of the acid and 
liberates it as a free gas. Instead of hydrochloric acid, we now have 
/inc chloride, which remains dissolved in the water in the generator. 
Zinc chloride is a white solid. 

All acids contain hydrogen, which may be replaced by certain 
metals. Hence, to prepare hydrogen, we may use almost any acid and 
one of a number of other metals instead of zinc. It is a curious fact 
that when pure zinc is added to pure hydrochloric acid, the chemical 

thistle tube 


Zn + HCI 


Fig. 6. Laboratory prepa- 
ration of hydrogen. Why 
does thistle tube extend 
below surface of liquid 
in the generator? 

hydrogen gas 
collecting jar 

- trough of water 



Fig. 7. Preparation of hydrogen 
by the action of sodium on water. 
Perform carefully. The reaction 
may be violent. 

action is very slow. However, slightly impure zinc replaces the 
hydrogen of the acid rapidly. The impurity in the zinc acts as a 

Can hydrogen also be prepared from water? Perhaps you have 
been thinking, "II the hydrogen of an acid can be replaced by a 
metal, can the hydrogen of water also be replaced by a metal?" The 
answer is Yes. Very active metals, such as sodium, potassium, and 
calcium, possess enough chemical energy to replace the hydrogen of 
water. This experiment may be tried by filling a bottle with water 
and inverting it in a dish partly filled with water, as shown in the 
illustration (Fig. 7) . Wrap in filter paper a piece of sodium the 
si/.e of a pea and put the paper in the coiled end of a wire. Then 
quickly insert this end of the wire into the bottle. Immediately, a 
very active evolution of gas is noticed and the bottle becomes filled 
with a colorless gas. The equation tor this reaction is: 

Sodium -f water sodium hydroxide -f hydrogen 
2Na +2HOH-> 2NaOH + H 2 

(Water may be represented as H 2 O or HOH.) 

In the chemical reaction that occurs, sodium replaces half of the 
hydrogen in water, and forms sodium hydroxide, NaOH. (The OH 
group is called the hydroxyl group.) The sodium hydroxide remains 
dissolved in the water. The water solution of sodium hydroxide feels 
soapy, and turns pink litmus, a vegetable coloring matter, blue. By 
evaporating the solution, the sodium hydroxide can be separated 
from the water as a white solid. Sodium hydroxide belongs to a group 
of compounds called bases (Chapter 14) . 

Even iron, which is not nearly as active as sodium, will replace the 
hydrogen of water while the water (as steam) is passed over red-hot 



iron in a heated tube. In tact, this method ol prcp.ning hvdrogen 
has been used to a slight extent, the products ol ihe ic.ution being 
iron oxide and hydrogen. 

What is the industrial method ol preparing hydrogen? In the 
order ol the quantities ot gas produced, hydrogen is obtained lot 
commercial use: (1) by passing steam oxer hot carbon, (2) by pass 
ing steam through natural gas (methane) , in the presence ol a 
catalyst, (3) by the electrolysis ol \\ater. The equations tor these 
methods are: 

1) Steam -f carbon > carbon monoxide -f hydrogen 

H '9 ! !+."".".". ?.: - CO + Hj 

2) Steam -f methane carbon monoxide -f hydrogen 

H>O + CH 4 -* CO + 3H, 

3 "i Water > hydrogen -f oxygen 
2HO-> 2H 2 4- O> 

In the first two methods, the hydrogen mav be sepaiated from 
tlie carbon monoxide gas by chilling the mixtuie ol gases. 1 lie cai- 
bon monoxide, whose free/ing point is much higher than that of 
hydrogen, solidifies, leaving nearly pure hydrogen gas behind. Only 
a minor part ol the total hydrogen production is by electrolysis. 

Physical properties of hydrogen. Hydrogen resembles oxygen in 
most ol its physical characteristics. It is a colorless, odorless gas, 
slightly soluble in water, and very difhcult to liquely. since it changes 
Irom a gas to a liquid at 252C. It differs physically from oxygen 
chiefly in its weight. It is the lightest element known. It is y 1 ^ as 
heavy as oxygen and ^^ as heavy as air. One liter ot hydrogen 
weighs approximately 0.09 gram at standard conditions' of tempera- 
ture and pressure. 

The metals palladium and platinum absoib large volumes ol 

hydrogen gas 
porous cup 

Fig. 8. Passage (diffusion) of 
hydrogen through a porous 
cup. Why is water forced 
through the glass tube? 

glass tube 



hydrogen. This absorption, or occlusion, is accompanied by such an 
increase in temperature that the metals actually glow. The absorp- 
tion of hydrogen by these two metals takes place in one type of 
lighting apparatus sometimes used in lighting the burners of gas 
stoves. Such a lighter usually consists of a fine wire or wires of one 
of the two metals strung between two suspension points. The lighter 
is placed above the burner. When the gas is turned on, hydrogen is 
absorbed by the wire. This causes the wire to give off heat. In about 
a second, the kindling temperature of the gas is reached, and the gas 
is ignited. 

If two or more gases are at the same temperature, the particles of 
the lighter gas move more rapidly than the particles of the heavier 
gases. Since hydrogen is the lightest gas known, its particles are in 
very rapid motion, passing through porous substances rapidly, as 
shown in Fig. 8. 

Hydrogen may burn quietly or explode violently. Although pure 
hydrogen burns quietly in air or in oxygen with a pale blue, almost 
colorless, flame, a mixture of hydrogen and oxygen may unite with 
explosive violence. The two gases, when mixed and kept below a 
temperature of about 800C., will not unite; but a spark, a flame, or 
a temperature of above 800C. will cause them to unite violently. For 
this reason, great care must be taken while experimenting with 
hydrogen to keep all flames away from the generator. It is also nec- 
essary to wait until the air has been completely expelled from the 
generator^ In -lore setting fire to the hydrogen as it escapes from the 
delivery wibe. , 

When the Frenchman, Pilatre de Rozier, heard of this gas which 
Cavendish had studied, he tried an unusual and foolish experiment. 

Fig. 9. Oxidation of hydrogen to form water. Will water form 
if bell jar becomes hot? How may bell jar be kept cool? 



He inhaled the gas until he had filled his lungs, and then as the 
gas issued from his mouth he set (ire to it. All Paris held its sides 
with laughter as it watched him spitting lire. However, when he set 
fire in the same way to a mixture of this gas and air. "the conse- 
quence was an explosion so dreadful that he imagined his teeth were 
all blown out." 

The chemical union of hydrogen and oxygen may be written: 

2H 2 -I- O 2 -> 2H 2 O 

This is an example of the strange behavior of chemical elements. 
Hydrogen, a highly flammable gas, unites with oxygen, a gas which 
helps things burn, forming water, a liquid that is one of the greatest 
enemies of fire. 

Experimental proof that water forms. Of course, when hydrogen 
combines with oxygen, we do not see the formation of a flood of 
water. We do not because the water that is formed at the tempera- 
ture of burning hydrogen is invisible since it is in the form of water 
vapor. However, it is possible to show the actual formation of water 
by arranging the apparatus shown in the illustration (Fig. 9) . As 
the invisible water vapor strikes the cool surface of the jar, drops of 
a liquid form. This liquid is pure synthetic water. A synthetic com- 
pound is a compound built up from simpler substances. 

Hydrogen is a powerful reducing agent. Since hydrogen has 
such a strong attraction, or affinity, for oxygen, it is able to tear 
oxygen away from the other elements of many of its compounds. 
The removal of oxygen from a compound is a process 
known as reduction. This ability of hydrogen to remove pxygen 
from the other elements of a compound makes it a reducing or 

deoxidizing agent. 

Fig. 10. Reduction of copper oxide by hydrogen. 

The fishtail burner is used to spread the flame, 
thus heating a larger surface of copper oxide. 
What is the function of the drying tube? Why is 
water a product? 

hydrogen generator 




For example, if pure hydrogen is passed over black copper oxide 
brought to a red heat, as shown in Fig. 10, the hydrogen takes the 
oxygen away from the copper oxide, leaving copper. This change 
may be represented as follows: 

Copper oxide (black) + hydrogen 
CuO 4- H 2 

> water 4- copper (red) 
H 2 O -f Cu 

Although hydrogen is a good reducing agent and is used to reduce 
the oxides of such metals as wolfram and molybdenum, other re- 
ducing agents are of greater commercial use. Perhaps the best 
example of these is carbon (coke) , which is used in reducing iron 
ore (iron oxide) to iron. 

The relation of oxidation to reduction. In the experiment just 1 
described, the copper oxide, CuO, is reduced to copper. At the same 
time hydrogen is oxidi/ed to water, H 2 O. This illustrates a general 
principle: namely, that the reducing agent is itself always oxidi/ed. 
You should remember that whenever one substance is reduced, 
another is oxidi/.ed. Thus oxidation and reduction always occur in 
the same reaction. 

The bunsen burner and how it works. One of the most familiar 
pieces of apparatus in the chemical laboratory is the bunsen burner 
(named after Robert Bunsen, a German chemist who introduced it in 
1855) . The function of this burner is to mix a gaseous fuel with 
air in order to make a lumliiminous flame that has a high tempera- 
ture and will not deposit soot. 

The parts of the bunsen burner designed for burning manufac- 
tured gas are shown in Fig! 1 1. Gas enters through the side of the 
stand, and its speed is increased by passing through the narrow spud. 
The rapidly moving gas draws in air through the hole in the collar, 
which may be adjusted to permit the correct volume of air to enter. 
The mixture of gas and air passes up through the barrel and is 
ignited at the top. 

When the pressure of the gas is low, the flame may be drawn back 
and may burn at the spud. This striking back of the flame can be 

Fig. II. 

of m 











dangerous because the stand -of the burner becomes hot. enough to 
produce serious burns when touched, and occasionally the hot burner 
melts the rubber hose and ignites the gas. 

In the type of burner for natural gas, the air intake is larger 
and the hole in the spud is much smaller. It is also supplied with 
a flame retainer which has small pilot jets led by gas bled from 
the main tube. This eliminates the tendency tor the flame to blow 

The burner on a gas range consists of a series of small bunsen 
burners. Not only is the nonluminous ilamc of the bunsen burner 
used in cooking, but it is used also to some extent in illumination by 
burning the gas inside gas mantles consisting of 09 pen cut thorium 
oxide and 1 percent cerium oxide. These oxides, when heated, glow 
with a rich white flame. They were first made commercially practical 
by von Welsbach (ton vels'baiO . The widespread use of electricity 
as an illimiinam has made this type of mantle burner all but obso- 

The nature of a flame. A flame is produced only when combustible 
vapors reach their kindling temperature in the presence of air, 
oxygen, or some other substance that supports combustion. When 
iron is heated gently, it glows, but a flame is not produced, because 
iron does not vapori/e at low temperatures. A candle, on the other 
hand, burns with a flame, because the candle melts and the heat 
from the burning wick causes this licjuid to change to a vapor that 
burns in air. It can be shown that the interior of a candle flame 
consists of a combustible gas by holding one end of a piping hot 
tube in the innermost part of the flame and then touching the other 
end with a lighted match (Fig. 12) . 

The structure of the bunsen flame. Three distinct /ones are no- 
ticeable in the bunsen burner flame as well as in the candle flame. 
The innermost zone A is composed of combustible gas that has not 
yet reached its kindling temperature. This fact may be tested by 
placing a match head in zone //. This may be clone by piercing 
a match near the head with a pin. The pin serves as a bridge across 

burning gas 

Fig. 12. Experiment thawing 
j the nature of the innermost 
: zone of a candle flame. 



the opening of the burner. If the burner is lighted carefully, the 
matcli will not catch fire. Outside zone A is zone B in which the gas 
is burning. The light-purple zone C is the region of complete com- 
bustion, where carbon dioxide and water vapor are formed. 

Just below the tip of the flame, plenty of air is available and, 
therefore, this is the oxidizing part of the flame. Zone B is the region 
of somewhat incomplete combustion, where one of the products 
formed is carbon monoxide, a reducing agent. Hence this zone is 
the reducing part of the flame. This zone of the flame is used in 
reducing metallic oxides to free metals by means of carbon. The 
blowpipe directs the flame, as shown in Fig. 71, page 321. 

The luminosity of the flame of the bunsen burner, when the collar 
is closed, is caused by the decomposition by heat of a small amount 
of the hydrocarbons and the subsequent burning of free carbon. This 
may be proved by holding a cold dish in the luminous flame. The 
carbon is reduced below its kindling temperature by contact with the 
dish and, hence, deposits soot on the dish. 

Hydrogen rides the winds. Interesting historically, but of only 
slight commercial importance today, is the use of hydrogen in filling 
balloons and other lighter-than-air craft. Soon after the discovery 
of hydrogen. Dr. Charles (shiirl) , of Paris, constructed the first 
large hydrogen-filled balloon, and in the presence of 300,000 spec- 
tators, Pilatre de Ro/ier, who had experimented unwisely with 
hydrogen before, bravely climbed inside its basket and started on 
the first aerial voyage ever made by a human being. Since that time, 
balloons filled with hydrogen have carried men around the world 
and have lifted explorers of the atmosphere to altitudes of almost 
14 miles. The development of the motor-propelled, rigid balloon 
called a dirigible gave us gigantic craft, nearly 1000 feet long and 
weighing more than 100 tons. 

From "Photographic History of the Civil War" by Albert Shaw 

The observation balloon 
was used to advantage by 
the Union Army in the War 
Between the States. Here 
the balloon Intrepid is being 
filled with hydrogen by the 
generators at the left. 

Such helium-filled bal- 
loons, carrying recording! 
instruments, are used to 
determine weather condi- 
tions at high altitudes. 

Official Unttcd States A'nrj/ J'/mf,, r ,-,,,,-, 

The history of this type of flying machine was filled with tragedy, 
to a large extent because hydrogen ignites and explodes readily. 
With the discovery of large sources of a nonflammable gas, helium, 
and its subsequent use, this ever-present danger was removed. There 
are no dirigibles in existence today, although blimps, which are 
motor-propelled, nonrigid balloons, are used for military purposes. 

Helium, which possesses about 93 percent of the lilting power of 
hydrogen, may be obtained today at a cost not much greater than 
that of hydrogen. The United States is fortunate in possessing the 
largest supply of helium of any nation. As high as seven percent 
helium is extracted from the natural gas of some western fields. 

To conserve our own supplies and prevent other nations from 
using helium for military purposes, Congress passed a law in 1938 
which placed its export under very strict government control. Sci- 
entific discoveries often turn out to be two-edged swords, and the 
careful control of such discoveries should be part of the business 
of government. 

Hydrogen, a gas, helps solidify fats and oils. One of the most 
important uses of hydrogen is based on the discovery that hydrogen 

Procter and Gamble Compan\ 

These machines, called 
"freezers," are used in the 
hydrogenation of vege- 
table oils to give the oils 
creaminess and smooth 



can be chemically combined with other substances to form new 
products of great value. This process of chemically combining hydro- 
gen with other substances is called hydrogenation. In normal times, 
hydrogenaiion ol fats and oils is the largest use of hydrogen. 

Generally, during hydrogenation, liquid oils are changed to semi- 
solid fats. Thus, cottonseed oil, in the presence of a finely divided 
nickel catalyst, unites with hydrogen, forming a white fat that is 
often sold nuclei such trademarks as "Crisco" or "Spry." Fish oils, 
formerly almost useless, have been hydrogenated similarly and ren- 
dered useful. Fats for soap-making and candle-making have been 
prepared commercially by this process also. Crude oil, coal, and cer- 
tain waste products of petroleum refining are also hydrogenated and 
changed in this way into high-quality gasoline. Wood alcohol is an- 
other important chemical made with the aid of hydrogen. 

Oleomargarine, or margarine, a widely used butter substitute, is 
made of either hydrogenated vegetable oils or animal fats, or mix- 
tures of these substances. Much of the margarine now marketed is of 
vegetable origin. 

The introduction of hydrogenation on a large scale helped the 
cotton farmers of the South by opening a new market for their cot- 
tonseed oil. It has had even greater effects abroad, for it has enabled 
several Kuropean countries to become less dependent upon other 
nations for their vital supplies of gasoline and edible fats. 

Hydrogen for the synthesis of ammonia. Ammonia, a compound 
of nitrogen and hydrogen, leads a double life. It is an important 
component of nearly all of the world's most powerful explosives, 
and it is an important ingredient in fertili/ers used all over the 
world. The most significant industrial process for the preparation of 
ammonia is one that embodies the direct combination, or synthesis, 
of nitrogen and hydrogen. This process is discussed in detail in 
Chapter 17. In wartime, production of ammonia probably requires 
more hydrogen than any other industrial use. 

In this plant, bituminous coal it 
hydrogenated to produce many 
useful chemicals. 

An illustration of a solar prominence. Solar prominences are great tongues of 
glowing hydrogen which shoot out of the chromosphere of the sun and extend 
far into space. The flames often attain lengths of more than 100,000 miles and 
have been known to reach lengths of more than 1,000,000 miles. 

Hydrogen in heating and cooling. Hydrogen has another impor- 
tant use. When it is burned, it gives off four times as much heat as 
an equal weight of coal. Because of the high temperature produced, 
hydrogen is used both in the oxy hydrogen torch and as a constituent 
of certain gaseous fuels such as water grw, which contains about 50 
percent hydrogen. Hydrogen gas is also used as a cooling agent in- 
stead of air. Because of its very low density it cuts down friction 
in machines such as high speed turbine-generators. When hydrogen 
is so used, the machine is completely enclosed to shut out air. 

The test for hydrogen. Hydrogen is not the only colorless gas that 
burns with a pale blue, almost invisible flame. But it is the only gas 
that forms water as the only product of its burning. This fact gives 
us a simple chemical test for hydrogen. 

Where is hydrogen found? Unlike oxygen, only small amounts of 
hydrogen occur on earth in the free state. However, live hydrogen 
is the most abundant element in the sun. The immense luminous 
tongues, or solar prominences, some of which extend half a million 
miles from the sun's surface, consist of glowing hydrogen. It is also 
the commonest element found in interstellar space, and by far the 
most abundant material out of which the whole universe is built. 

Combined hydrogen is very common on earth. Hydrogen consti- 
tutes about 11 percent by weight of all water and is one of the ele- 
ments in petroleum, all acids, and living cells (protoplasm). In 
spite of its widespread occurrence, the extreme lightness of hydrogen 
accounts for the fact that it constitutes only 1 percent by weight of 
the earth. 


Henry Cavendish (1731-1810), one of 
the most unusual personalities in the 
history of science. For the most part, he 
shunned society of any kind, even in- 
structing his servants to keep out of his 
sight. Yet, over his long lifetime, he 
made many important contributions in 
the fields of chemistry and physics. 

Theories lead to great discoveries. In 1932, three Americans 
headed by Harold C. Urey discovered that ordinary hydrogen could 
be separated into two distinct forms. They named the heavier of 
these forms deuterium. Two years later, it was proved that there 
is a third form of hydrogen. This form, the heaviest of the three, 
has been named tritium. The three forms have the same chemical 
properties, but differ in certain physical properties. 

This achievement is remarkable, not so much because ordinary 
hydrogen was shown to be a mixture of three forms of the same 
element, but because it afforded a definite example of a great con- 
tribution to chemistry made by scientists who forecast these dis- 
coveries purely from theory. Their success points to the fact that 
science needs the man who experiments, the thinker who can put 
his imagination and reason to work in propounding theories, and the 
engineer who works to discover how these theories and processes 
may be used in the service of man. 

Another example of how theoretical problems in science may turn 
out to be of great practical value, was the theorizing of the great 
mathematician, James Clerk Maxwell. In 1863 he came to the con- 
clusion that just as light results from a wave disturbance in the 
ether, so electric disturbances from a spark should produce similar 
waves, invisible, to be sure, but nevertheless existent. Experimental 
evidence of such waves was found 23 years later by a young physicist, 
Heinrich Hert/. These Hertzian waves, now known as radio waves, 
were later used by Marconi in the transmission of wireless messages. 
Thus modern radio originated, and out of a purely theoretical in- 
vestigation came one of the most practical and valuable of modern 
scientific marvels. Often, even the greatest scientist cannot predict 
the practical value of theoretical research. 





Hoyle, Fred. The Nature of the Universe. Harper and 
Brothers, New York, 1950. This is a very small book that gives 
an exciting picture of the composition of the stars and inter- 
stellar space. 

Jaffe, Bernard. Men of Science in America, pp. 3S9-355. 
Simon & Schuster, New York, 1944. Interesting inioimation on 
early American work in aeronautics. 

Walters, Leslie. "Chemistry Exhibits and Projects." Jour- 
nal of Chemical Education, March, 1939, pp. 113-115. An 
illustrated article on exhibits and projects made by high- 
school students. Suggestions for what you, too, could do along 
this line. 


1. A reducing agent removes oxygen from a compound con- 
taining oxygen. (This definition will be somewhat expanded 

2. Reduction and oxidation always occur in the- s.nnc re- 

3. Occlusion is the absorption of gases by metals. 

4. Careful control of some scientific discoveries is necessary 
to prevent their misuse. 

5. Chemistry needs the experimental chemist, the scientist 
able to use creative imagination in formulating theories, and 
the engineer who works to discover how theories and processes 
may be used in the service of man. 


Group A 

1. (a) Who is credited with the discovery of pure H 2 ? 
(b) How was H 2 first prepared? 

2. (a) What substances are used in the laboratory prep- 
aration of H 2 ? (b) How do we know that the H.. comes iroin 
the acid and not from the metal? (c) Write the word-equation 
for this reaction. 

3. (a) What advantage is there in collecting H 2 by the 
displacement of H 2 O? (b) What kind of gas could not be 
collected in this manner? 


4. (a) Make a labeled diagram of the laboratory prepara- 
tion of H 2 . (b) Why should the thistle tube extend below the 
surface of the liquid in the generating bottle? (c) What are 
two reasons for using a thistle tube? 

5. (a) When Na displaces H 2 from H 2 O, what is the other 
substance formed? (b) Why can you not see it? (c) How can 
you obtain it for inspection of its properties? 

6. What metals other than Na are so active that they will 
liberate H 2 from water when the metal is simply placed on 
H 2 O? 

7. (a) What metal will liberate H 2 from H 2 O under cer- 
tain special conditions? (b) Describe this method of prepar- 
ing H 2 . 

8. (a) In what two ways is H L , prepared tor commercial 
use? (b) Why were these methods selected in preference to 
ones used in the laboratory? 

9. What precaution must be observed in preparing and 
handling H 2 ? 

10. (a) List five physical properties of H 2 . (b) Describe a 
simple experiment to illustrate the tact that H 2 is a very light 

11. Determine the weight and cost oi the H 2 that would be 
needed to fill a dirigible of seven-million-cubic-foot capacity. 
Consider that the cost of the H 2 is $2.00 per hundred cubic 
feet of gas. 

12. What element in the air is used in the burning of H 2 ? 


13. How can you show that water is formed when hydrogen 

14. What is reduction? 

15. In the reduction of CuO by H 2 , what substance is oxi- 

16. How can you identify each of the products obtained in 
the reduction of CuO by H 2 ? 

17. Oxidation and reduction always occur together. Ex- 

18. (a) Give an example of reduction carried out on a large 
scale in industry, (b) What reducing agent is used? 


19. (a) What is the function of each part of a bunsen 
burner? (b) What is meant by the striking back of the flame? 

20. (a) What conditions are necessary for a flame? (b) 
Make a labeled drawing of the flame of a bunsen burner. 

21. What formerly useless byproduct is converted into a 
very useful substance by the use of H ? 

22. How is oleomargarine made? 

23. In what special fuels is H 2 chiefly responsible for the 
high temperatures obtained? 

Group B 

24. Can you suggest a sale way to test H, for its burnability 
as it emerges from the end of the delivery tube in the labora- 
tory preparation? 

25. Helium gas is twice as dense as hydrogen gas and yet 
can lift about 93% as much weight as hydrogen gas. Explain. 

26. (a) Could pure H 2 be used in the gas range at home? 
(b) Explain. 

27. Why does pure H 2 burn quietly in an atmosphere of air, 
yet burn with explosive violence when the two gases are mixed 
and ignited? 

28. Tightly bound inflated balloons gradually collapse. Ex- 


1. A student's mother stopped using "Crisco" after her son 
came home with the news that "Crisco" was not a "natural 
product" but was manufactured chemically, (a) Was the 
mother justified? (b) Give reasons for your answer after 
you have consulted your family physician, the producers of 
"Crisco," and the U.S. Department of Agriculture. 

2. Organize a small discussion group in your chemistry class 
and discuss the topic, "Synthetic chemistry has helped in the 
rise of totalitarian states." 

3. Prepare a report or organi/e a class discussion on the 
topic "Theoretical science has irequently resulted in great 
practical discoveries." 

4. WATER: 


. . . Laboratories are necessary, and, 
though an artist without a studio or 
an evangelist without a church might 
conceivably find under the blue 
dome of heaven a substitute, a sci- 
entific man without a laboratory is a 
misnomer. Frederick Soddy, 1877- 

Water a compound of two gases impossible! For thousands of 
years, water was considered an element. Aristotle, one of the wisest 
of Greeks, included this liquid among the "four elements" of the 
ancients. No power of man seemed strong enough to break it up 
into any recognizable components. Although it is true that by 1780 
a number of scientists had really decomposed water so that hydrogen 
was liberated, they were unaware of the real nature of what they 
had done. They could not believe that hydrogen had actually come 
from the water. 

"It is very extraordinary that this fact should have hitherto been 
overlooked by chemists. Indeed, it strongly proves that in chemistry 
it is extremely difficult to overcome prejudices imbibed in early 
education." These were the words of Lavoisier, Often it is hard to 
overcome superstition, prejudice, and tradition. However, we must 
learn to do so in order to think and act scientifically. 

In 1784, Cavendish, who had studied hydrogen, read an exciting 
paper before the members of the Royal Society of England. This 
is what he told them: "Water is a compound of oxygen and hydro- 
gen." What a startling announcement! Water a compound of two 
colorless, tasteless gases? What were his proofs? Cavendish told them 




quietly and without emotion. He said that he had placed in a glass 
flask a mixture composed of about twice as much air as hydrogen. 
Then he had passed an electric spark through the mixture. "All the 
hydrogen and about one-fifth of the air condensed into a dew which 
lined the glass. In short," he continued, "it seemed pure water." His 
experiments proved conclusively that water is a compound of oxygen 
and hydrogen, and yet Cavendish said "it seemed." He suspected his 
listeners would not be convinced. Water a compound of two gases 

Lavoisier convinces the world that water is a compound. Lavoisier, 
who had explained the nature of burning, determined to tear apart, 
or analyze, water by an experiment that would convince the world, 
just as Cavendish had shown the world that he had built up, or 
synthesized, water from oxygen and hydrogen. 

He arranged the apparatus shown below. In the retort A he heated 
pure water, so that steam would pass through the tube containing 
pure charcoal, which, as you see, was heated in a furnace. The gas 
that escaped passed into the jar H. He found carbon dioxide gas 
dissolved in the water in the jar and identified the gas issuing from 
H as hydrogen. The water (steam) had been broken up into hydro- 
gen, which passed on as a gas, and oxygen, which combined with the 
carbon, forming carbon dioxide. Lavoisier collected and weighed 
both gases. The weight of all the resulting products accounted for 
all changes in the weights of the substances used in the reaction. 
Thus Lavoisier proved without a doubt that water could be broken 
up into the two gases of which it is composed hydrogen and 

Fig. 13. Apparatus used by Lavoisier to 
analyze water. Compare this equipment with 
that used in a modern laboratory. 

burning charcoal 


hydrogen gas 

reservoir of water containing 
some H 2 SO 4 


Fig. 14. Electrolysis of water. Why is a reservoir necessary in this apparatus? 

The equation for the chemical changes that took place in this 
experiment is: 

2H 2 O + C -> 2H 2 + CO 2 (carbon dioxide) 

How water can be broken up by an electric current. In 1800, the 
invention of the electric battery by Volta put into the hands of 
scientists a new and powerful tool for determining the composition 
of compounds. The decomposition of any compound by electricity 
is called electrolysis. 

Within a few months following Volta's invention, an apparatus 
had been devised for the electrolysis of water. This special piece of 
apparatus is used in the laboratory in determining the composition 
of water by volume. Water is poured into the center glass tube, 
which acts as a reservoir, and fills the other two tubes, as shown 
in Fig. 14. Since pure water is an extremely poor conductor of 
electricity, that is, will not allow electricity to pass through it readily, 
a small amount of sulfuric acid is added. (At the close of the chem- 
ical action, the amount of sulfuric acid is unchanged.) 

The electric current enters the apparatus through one of the 
platinum electrodes, passes through the mixture of water and sulfuric 
acid, and leaves the apparatus at the second platinum electrode. 
The source of current may be any source of direct current, such as 
a storage battery or dry cells. As soon as the circuit is closed, bubbles 
of gas collect on the surface of the electrodes, rise through the water, 
and collect at the top of the two outside tubes. The gas that collects 
at the negative electrode, or cathode, is hydrogen; the gas that col- 
lects at the positive electrode, or anode, is oxygen. 



The composition of water by volume. No matter when we stop 
the action of the current in the electrolysis of water, we discover a 
singular fact. The volume of the gas at the cathode is twice the 
volume of the gas at the anode. In other words, water always consists 
of 2 parts of hydrogen to 1 part of oxygen by volume. 

If the gases hydrogen and oxygen are mixed in a closed chamber 
and then exploded by an electric spark, water vapor is formed, and 
the ratio in which these gases unite, forming water vapor, is always 
2 volumes of hydrogen to 1 of oxygen. Any excess of either gas is left 

The composition of water by weight. If pure water is thus de- 
composed, forming two gases in a definite ratio by volume (vol- 
umetric ratio) , and if water is formed by combining these gases in 
the same volumetric ratio, it ought to be possible to find in this com- 
bination a constant ratio by weight. This is what Cavendish did in 
1784 when he formed water by combining these two gases. He actu- 
ally weighed both the gases and the water. Today we know that if 
we combine 1 gram of hydrogen (11 liters) with 8 grams of oxygen 
(5.5 liters) , we obtain 9 grams of water. Thus, the composition of 
water by weight is 1 part of hydrogen to 8 parts of oxygen. 

The composition of water by weight may be tested in the labora- 
tory by making use of the fact that pure hydrogen passed over heated 
copper oxide reduces the copper oxide, yielding free copper, by 
uniting with the oxygen of the copper oxide, thus forming water 
(see page 50) . The weight of the water formed is always found 
to be exactly the same as the weight of the hydrogen used plus the 
weight of the oxygen lost by the copper oxide. The following data 
are the results of such an experiment: 

Weight of copper oxide before experiment 80 grams 

Weight of copper left after experiment 64 grams 

Weight of oxygen that combined with hydrogen 16 grams 

Weight of water produced 18 grams 

Weight of hydrogen used 2 grams 

Result: 16 parts by weight of oxygen combined with two parts of 
hydrogen (or 8 parts of oxygen combined with one part of hydro- 
gen) , forming water. 

. induction 

latmum electrodes co ;| 

Fig. 15. Eudiometer. In this appara- 
tus, a mixture of hydrogen and oxy- 
gen is ignited by an electric spark, 
forming water. Any excess of either 
gas acts as a cushion against the 
rush of mercury up the tube. 

Accuracy and satisfaction in scientific investigations. An exact 
determination of the relative combining weights of oxygen and hy- 
drogen was essential to the mathematics of chemistry. Hundreds of 
men performed thousands of experiments to determine these values 
as accurately as human ingenuity could devise. Many eminent Ameri- 
cans were among them. Edward Morley (1838-1923) , a professor at 
Western Reserve University, spent more than ten years of his life on 
such experiments. Finally he arrived at a number which still stands 
as the basis for chemical calculations. He found that 1.008 parts of 
hydrogen by weight combine with eight parts of oxygen by weight, 
forming water. 

Morley and the others did not receive monetary rewards from 
these experiments. Theirs was a labor of love, a work of pure scien- 
tific research. Their sole reward was the satisfaction of knowing that 
they were helping to increase scientific knowledge whjch might be 
of use to all humanity. The world needs such unselfish men. 

The law of definite proportions. You have learned that the com- 
position of water is always the same. Perhaps you have wondered 
if this is a property of all chemical compounds. It is one of the 
fundamental laws of chemistry that, in forming compounds, elements 
combine in exact proportions. In fact, constancy of composition is 
the most valid test used in deciding whether a substance is a com- 
pound or a mixture. In other words, the composition of a pure com- 
pound never varies. This is the law of definite proportions.. 

"The stones and soil beneath our feet and the ponderous moun- 
tains are not mere confused masses of matter; they are pervaded 
throughout their innermost constitution by the harmony of num- 
bers." This is indeed fortunate, for if the composition of pure com- 
pounds ever changed, instead of being always the same, the exact 
measurements that we use in chemistry would not be possible. Par- 
ticularly is this true of quantitative chemistry, for if the composition 
of compounds varied, we should have no reliable standard chem- 
icals. Standards are necessary, not only as a means of judging the 
purity of a substance which we plan to use, but also as a reliable 
basis for comparison. 



What is a "pure" chemical? Absolutely pure chemicals are almost 
impossible to make or buy. Chemicals that have no appreciable trace 
of impurities are called C.P., that is, chemically pure. The designa- 
tion U.S.P. refers to standards of purity of chemicals to be used in 
medicines listed in the United States Pharmacopoeia. Such chem- 
icals contain no harmful impurities. Reagent grade chemicals con- 
form to the standards of the American Chemical Society. Tech chem- 
icals do not meet any definite, or fixed, standards of purity but are 
suitable for many uses in which slight impurities are of little im- 

Patent medicines, other remedies, and drugs, such as aspirin, are 
frequently advertised and sold under trademarks. The purchaser 
should always carefully examine the label on the package in order 
to be sure of the purity of the contents. The chemical composition 
of the contents will also serve as a guide to the fair price of the 
article. Occasionally, simple and inexpensive C.P. chemicals, such 
as bicarbonate of soda, are masked behind trademarks and sold at 
exorbitant prices. The consumer can help to guard against such 
practices by insisting that accurate and complete statements of com- 
position be printed on labels of all packaged goods. 

Some physical properties of water. As you know, water is an odor- 
less, tasteless liquid that is colorless, except in very thick layers, when 
it appears blue. Pure water freezes at 0C. or 32F. and boils at 
100C. or 212F., at standard conditions. In general, impure water 
has a higher boiling point and a lower freezing point than pure 

The fact that water is so universally distributed has led to its use 
in the devising of scientific standards of measure. Thus a gram, the 
metric unit of weight, is, by definition, the weight of a milliliter * of 
chemically pure water at 4C. This temperature is chosen because, 
when water cools, it contracts until the temperature reaches 4C. 
Below that temperature it begins to expand again. Hence, 4C. is 
the temperature at which a unit volume of water weighs the most. 
The weight of a unit volume of a substance* is known as its density. 
Thus at 4C. water has its greatest density. Since we use the gram 
and the cubic centimeter or the milliliter as our units of weight and 
volume, we may redefine the density of a substance as the weight in 
grams of 1 cubic centimeter or 1 milliliter of that substance. 

* A milliliter (ml.) is a small unit of capacity (volume) in the Metric System. 
Numerically it is equal to 0.001 of a liter and is used in measuring the volume of 
fluids. The cubic centimeter (cc.) is also used in measuring volumes, particularly 
of solids. 







water 1.0 

Fig. 16. The relationship of specific gravity and 
buoyancy. Aluminum has a specific gravity greater 
than that of water and does not float. Ice has a specific 
gravity slightly less than that of water and floats 
largely submerged. Cork has a low specific gravity 
and floats with most of its mass above water. 

Since water has a density of 7, that is, 1 milliliter of water weighs 
1 gram at 4., the density of any substance is also the ratio of the 
weight of a given volume of that substance to the weight of an equal 
volume of water. We call this ratio the specific gravity (sp. gr.) of 
the substance. It shows the comparison between the weight of the 
substance and the weight of an equal volume of water. For example, 
when we say that concentrated su If uric acid has a specific gravity 
of 1.84, we mean that it is 1.84 times as heavy as water, volume for 
volume. Since below 4C. water expands, ice is lighter than water 
and floats on it. Ice, of course, has a specific gravity less than that of 
w r atcr. 




(In grams per cubic centimeter or per milliliter) 

0.97 I Iron (pure) 7.86 | Lead 



19.3 I Platinum 



The high specific heat of water. The amount of heat necessary to 
raise the temperature of 1 gram of water 1 degree centigrade is called 
a calory. The number of calories necessary to raise the tempera- 
ture of 1 gram of a substance 1 degree centigrade is known as the 
specific heat (sp. lit.) of that substance. Because it takes 1 calory to 
raise the temperature of 1 gram of water 1 degree centigrade, the 
specific heat of water is 1. 

Since it takes only one-thirtieth as much heat to raise the tempera- 
ture of 1 gram of mercury 1 degree centigrade as it does to raise 1 
gram of water 1 degree centigrade, the specific heat of mercury is one- 
thirtieth of 1, or^j. 

Water has a higher specific heat than most other substances. Since 
it requires so much heat to raise its temperature, it warms up slowly. 
Conversely, upon cooling it gives up a larger amount of heat for the 
same fall in temperature than most other substances do. It is partly 
because of the high specific heat of water that it is used in home heat- 
ing systems and in the cooling systems of automobiles.* 

The chemical properties of water. Water is a stable compound, 
that is, it cannot be decomposed easily. It does not even begin to 
decompose into hydrogen and oxygen until a temperature of 
1000C. is reached. Even at 2500C. only two percent of it is decom- 
posed. However, electricity, in the presence of a catalyst, tears it 
apart easily (see page 62) . 

At ordinary temperatures, water is decomposed by the more active 
metals, such as sodium and potassium, and at higher temperatures 
by the less active metals, such as iron. In these cases the gas liberated 
is hydrogen. Water is also decomposed by the more active nonmetals, 
such as chlorine and bromine, but these liberate oxygen from water 
instead of hydrogen. 

Water acts as a catalyst in many chemical reactions. For example, 
perfectly dry oxygen and hydrogen do not unite when a spark is 
passed through them, yet the faintest trace of water causes such a 
mixture to explode. Phosphorus does not burn in perfectly dry air, 
but burns readily if even a trace of water vapor is present. 

* Although the temperatures of boiling water and steam are the same, it takes 
about 540 calories to change 1 gram of water from its boiling point of 100C. 
to steam at 100C. This amount of heat is called the heat of vaporization of 
water. Real steam, or water vapor, the gaseous form of water, is invisible. The 
visible cloud commonly called steam is water vapor after it has condensed into 
tiny liquid droplets. 

In ice water the temperatures of the freezing water and the melting ice are 
both 0C. Yet it requires about 80 calories to change 1 gram of ice at 0C. to 
1 gram of water at 0C. This amount of heat is called the heat of fusion of ice. 

(left) Photomicrograph of a crystal of sodium carbonate, (right) The same crys- 
tal after a few hours of exposure to air. What has occurred? 

Water of crystallization. A crystal is a solid mass having a well- 
clefined and angular form. The word is derived from a Greek word 
meaning clear ice. Most elements and compounds are capable of as- 
suming the crystalline form, showing sharp edges and flat surfaces. 
Such a substance is crystalline washing soda, or sodium carbonate. 

When a crystal of washing soda is heated or even exposed to air, 
it gives off water and crumbles to a white powder which is not 
crystalline. The weight of water liberated bears a fixed ratio to the 
weight of the crystal and is united chemically with the compound of 
which the crystal is composed. Water which is thus chemically united 
with a substance and gives that substance its crystalline form is called 
ivater of crystallization. Such water is rather loosely held in chemical 
combination and may be easily expelled. The water of Crystallization 
is separated from the rest of the formula by a centered dot, which 
means plus (+) and is not a multiplication sign. A substance that 
contains water of crystallization is sometimes called a hydrate. 

Another common hydrate is crystallized copper sulfate once com- 
monly known as blue vitriol. When this compound is heated, its 
water of crystallization is liberated and it crumbles to a white powder. 

CuS0 4 5H 2 -> CuS0 4 + 5H 2 
crystallized copper sulfate water of 

copper sulfate (anhydrous) crystallization 

This change in color is further evidence that the water of crystalliza- 
tion is chemically united with the copper sulfate. Use is made of the 
difference in color between white anhydrous copper sulfate and the 
blue hydrated copper sulfate as a test for water. Water will change 
anhydrous copper sulfate to the blue hydrate. 

The ability to form crystals is not always dependent upon the pres- 
ence of water. Many crystalline substances, such as table salt (so- 
dium chloride) arid sugar, do not contain water of crystallization. 
They are said to be anhydrous, meaning without water. Crystals that 
have lost their water of crystallization are also said to be anhydrous. 




Efflorescent substances give up water. Crystallized washing soda, 
on exposure to air, loses its water of crystallization and crumbles 
to a powder. Such a substance is said to be efflorescent, which means 
that it gives up its water of crystallization on exposure to air. The 
drier the air, the faster the loss of water of crystallization. 

Deliquescent substances take up water. Dry sodium hydroxide, 
when left exposed to air, soon absorbs enough water from the at- 
mosphere to dissolve itself in this water. Such a substance is said to 
be deliquescent. The higher the percentage of water vapor in the 
air, the faster the process of deliquescence. 

Calcium chloride, a white solid, is deliquescent and is often used 
to sprinkle dry roads and tennis courts. It absorbs moisture from the 
air and, in this way, helps to keep the dust down. Magnesium chlo- 
ride, an impurity found in common table salt, is also deliquescent. 
Removal of the magnesium chloride causes pure table salt to remain 
dry in damp weather, to pour easily, and not to cake. (What sub- 
sance does your mother put in a salt shaker to keep the salt from 
becoming lumpy?) 

Deliquescent substances may be used as drying, or dehydrating, 
agents. For example, concentrated sulfuric acid absorbs moisture 
from the air and, therefore, is used in drying gases. When used in 
the laboratory, these dehydrating agents (the most common of which 
are sodium hydroxide, sulfuric acid, and calcium chloride) are often 
placed in the lower compartment of a vessel known as a desiccator; 
the upper compartment, only partially separated from the lower, 
contains the substance to be dried. 

Our lives depend on water. Water is essential to life. Almost 70 
percent of the total weight of the human body is water, and plants 
contain even more. Lettuce, for example, contains as much as 95 
percent water by weight. 

,,.. .:- *v ! -'" - Fig. IF* Desiccator. 5yb 

i. * to be are 

l n the VGpw m4 the 

is put In 



Immense tracts of land in our own country, such as the hitherto 
arid Columbia Basin of the Northwest, have been or will be turned 
into rich farmland by irrigation. Federal government projects have 
included construction of huge dams such as the Norris Dam, the 
Hoover Dam, the Fort Peck Dam, and the Grand Coulee Dam. Be- 
hind mountainous walls of earth and concrete are stored huge res- 
ervoirs of water which are changing more wasteland into fertile 
plains and are helping to solve the problem of the frequent recur- 
rence of disastrous floods in certain areas. 

Water is found in rocks, paper, fibers, and other substances gen- 
erally thought of as "dry." The pages of this book may contain as 
much as 10 percent water by weight. The importance of water, the 
most common solvent in nature, is discussed in detail in Chapter 15. 

Water and health. Water is a major component of all body tissues 
and fluids. It plays a major role in the preparation of foods we eat 
and in the processes of digestion and assimilation. Both nutrients and 
oxygen are carried to the cells of the body in fluids composed chiefly 
of water, and many of the waste products of the body are carried 
away and eliminated in a similar mariner. \Vater, in the form of 
perspiration, aids in regulating the temperature of the body. 

To maintain normal body processes, rather large quantities of 
water are necessary. The amount of water a person ^should drink 
each day to enable these processes to be carried on varies with the 
kind and amount of activity, the temperature, and various other 
factors. However, 6 glasses each day should be considered a minimum 
for good health. 

The idea that water should not be drunk with meals is without 
foundation. Digestion proceeds normally even when large quantities 
of water are present in the stomach. However, it is very important 
not to substitute the drinking of water with meals for proper and 
complete chewing. 

This illustration shows how a single calcium chloride crystal deliquesces, gradually 
absorbing water from the air until it becomes completely dissolved. 




Holmes, Harry N. Out of the Test Tube (4th ed.) , 
Chap. VII. Emerson Books, New York, 1945. "The Elixir of 
Life," of course, refers to water. 

Read the Label on Foods, Drugs, Devices, Cosmetics. Cata- 
logue No. FS 13.111:3/2, 1953. Supt. of Documents, Govt. 
Printing Office, Washington, D.C. Revision No. 1 of a 35-page 
illustrated pamphlet containing valuable information. 15^ 

Thorpe, T. E. Essays on Historical Chemistry, pp. 98-122. 
The Macmillan Co., London, 1923. A description of the 
famous controversy (the Water Controversy) over the priority 
of the discovery of the composition of water, involving Cav- 
endish, Watt, Priestley, and Lavoisier. 


1. Analysis is the breaking down of a compound into 
simpler substances. 

2. Synthesis is the building up of a more complex com- 
pound from simpler substances. 

3. The decomposition of any compound by electricity is 
called electrolysis. 

4. Edward Morley spent more than ten years of his life in 
determining the exact ratio in which oxygen and hydrogen 
unite in forming water. His work is a fine example of ac- 
curacy and patience in scientific research. 

5. Every pure chemical compound has a definite composi- 
tion. This is the law of definite proportions. C.P. refers to a 
chemically pure substance. 

6. The density of any substance is the weight of a unit 
volume of that substance. 

7. A calory is the amount of heat necessary to raise the 
temperature of one gram ol water one degree centigrade. 

8. The specific heat of a substance is the number of calories 
necessary to raise the temperature of one gram of that sub- 
stance one degree centigrade. 

9. The specific gravity of any substance is the weight of 
one cubic centimeter or one milliliter of that substance com- 
pared with the weight of an equal volume of water. 

10. Water of crystallization is the water chemically present 
in certain crystalline substances. 

11. An efflorescent substance loses water of crystallization 
on exposure to air; a deliquescent substance absorbs water 
from the air. 



Group A 

1. Why is water considered a compound? 

2. How did Cavendish synthesize H 2 O? 

3. What is the difference between analysis and synthesis? 

4. Briefly describe how Lavoisier analyzed H 2 O. 

5. What is electrolysis? 

6. (a) Make a labeled diagram of the apparatus for the 
laboratory electrolysis of H 2 O. Indicate the direction of the 
current, the cathode, and the anode, (b) At which electrode is 
the H 2 given off? (c) the O 2 ? 

7. What part does sulfuric acid (H 2 SO 4 ) play in the elec- 
trolysis of H 2 O? 

8. Write the word-equation for the electrolysis of H 2 O. 

9. How would you test to find out which of the gases 
present in the two outer tubes of the electrolysis apparatus 
is H 2 ? 

10. (a) What is the composition of H 2 O by volume? (b) by 

11. What are the differences between C.P., V.S.P., Tech, 
and Reagent chemicals? 

12. In connection with the study of H 2 O, cite an example of 
the accuracy and patience ol men of science. 

13. State and illustrate the law of definite proportions. 

14. (a) Does vigorously boiling H 2 O have a higher tempera- 
ture than slowly boiling H 2 O? (b) Explain. 

15. What are five physical properties of H 2 O? 

16. How does steam differ from water vapor? 

17. Why is it sometimes unsafe to purchase drugs or medi- 
cines by trade-mark alone? 

18. Explain why water pipes often burst in cold weather. 

19. (a) Which would you prefer to heat your feet on a cold 
night a hot flatiron or a bottle of hot H 2 O? (b) Explain. 

20. If H 2 O were less stable than it is, what disaster might 


21. Describe two chemical properties of H O. 

22. Give an example of the part that a trace of H 2 O may 
play in bringing about a chemical change. 

23. Illustrate what is meant by water of crystallization. 

24. Transparent crystals of washing soda change to a white 
powder. Is this a physical or a chemical change? 

25. Why is an unstoppered bottle of calcium chloride 
(CaCl 2 ) sometimes left in large clocks? 

26. When crystals of table salt (NaCl) are heated, some 
H 2 O is liberated. Is this water of crystallization? Explain. 

27. Examine some pellets of NaOH that have been exposed 
to air. What property does NaOH have? 

28. (a) Which gives a more severe burn, boiling H 2 O or 
steam? (b) Why? 

29. Explain the importance ol H 2 O to health. 

Group B 

30. If H 2 O did not expand on Iree/ing, how would it affect 

31. A spark is passed through a mixture ol 60 ml. of O 2 and 
50 ml. of H 2 in the presence of water vapor. What substances, 
and how much of each, will be found in the tube after the 

32. Devise an experiment to show the composition of H.,O 
by weight. Mention actual weights. 

33. How would you determine the percentage of H.,O 
present in a sample of "dry" wood? 

34. A liter of H weighs 0.09 g. and a liter of O 2 weighs 
1.43 g. Show how you could find the composition of H 2 O by 
weight from these facts (data) . 

35. What weight of oxygen can be obtained from the elec- 
trolysis of 50 pounds of water? 

36. Ice is purer than water. Would it be safe to use ice from 
a polluted pond in your iced tea? Explain. 


1. Find out from your mother or grocer the cost of pound 
packages of powdered *washing soda and crystallized washing 


soda. The water of crystallization present in crystallized wash- 
ing soda is equal to approximately 63 percent of the weight 
of the crystal. On this basis, calculate which is less expensive, 
the crystallized washing soda containing water of crystallization 
or the powdered washing soda containing no water of crystal- 
lization. Make a report of your problem in class. 

2. A student came to his chemistry teacher very excited 
about his invention of a system that would run an airplane 
indefinitely on a small amount of H 2 O. He planned to de- 
compose H 2 O by electricity, use the O 2 and H 2 thus produced 
to supply an oxyhydrogen torch which, when ignited, would 
boil H 2 O for a steam engine. The steam engine in operation 
would give power to the plane's propeller and, at the same 
time, to an electric generator. The electricity thus produced 
would be harnessed to decompose more H 2 O, which would be 
constantly renewed as a product of the burning of the O 2 and 
H 2 in the oxyhydrogen torch. What do you think of this in- 
vention? Your answer should be complete with scientific 

3. Benjamin Thompson, one of the greatest of early Amer- 
ican scientists and who later became Count Rumford, made 
classic contributions in the field of specific heat. "It is a strange 
coincidence," wrote Albert Einstein, "that nearly all the funda- 
mental work concerned with the nature of heat was done not 
by professional scientists but by men who regarded science as 
their great hobby." Mention two other scientific contributions 
made by men who were not professional scientists. 

4. Get some anhydrous copper sulfate and with it determine 
whether certain "dry substances" really contain water. Report 
your results. 


The bodies which time and nature 
add to things little by little, . . . 
no exertion of the eyesight can be- 
hold; and so, too, wherever things 
grow old by age and decay, and when 
rocks hanging over the sea are eaten 
away by the gnawing salt spray, you 
cannot see what they lose at any 
given moment. Nature, therefore, 
works by unseen bodies. 

Lucretius, 99-55 B.C. 

The origin of the idea of atoms. The most fruitful scientific specu- 
lation that came out of ancient Greece 2500 years ago was that mat- 
ter is made up of small, eternal particles in continual motion. Leu- 
cippus and his pupil Democritus (de mok'ri tws) , taught that all 
matter was composed of invisible, indivisible, indestructible parti- 
cles, or atoms. During the seventeenth century, Newton had similar 
ideas about the nature of matter. "It seems probable to me," wrote 
Newton, "that God in the beginning formed matter in solid, massy, 
hard, impenetrable, moveable particles, so very hard as never to wear 
or break into pieces." The gradual development of the idea of the 
atom is an interesting story. 

How Dalton's approach to the nature of matter differed from that 
of the Greeks. Early in the nineteenth century, an English scientist, 
John Dalton, became interested in the idea of atoms. Newton's idea 
was a beautiful one, thought Dalton, but did it check with the known 
facts? Would it help explain some of the physical properties of gases, 
which had so puzzled him? 

It is interesting to note the difference in the way the Greek teach- 
ers had come to their ideas and the way in which Dalton formulated 
his theory. The Greek teachers made a few observations, followed. 




some logical reasoning, and then ventured an opinion. For example, 
they noticed that a lump of salt could be broken down into bits of 
salt, which could then be further reduced to tiny crystals of salt. It 
was inconceivable, they reasoned, to continue this division forever. 
There must come a time, therefore, when one would finally reach the 
smallest piece of salt, that is, an indivisible atom of salt. Dal ton did 
more. He also experimented. He tried to find out why the gases of 
the atmosphere remained mixed, how gases dissolve in water, and 
whether the composition of pure compounds varies or is constant. 
On the basis of some not too accurate observations in these experi- 
ments and some indirect evidence from facts known in his day, 
he formulated the atomic theory. In 1803 he announced it suddenly 
without waiting to test all of it by experiment. Since that time a 
great many additional experiments have proved that Dalton's theory 
was substantially correct. , ; 

With this electron microscope, 
scientists may study molecules 
and other infinitesimal par- 
ticles. Typical photomicro- 
graphs may be seen in the 


Dalton's atomic theory. Dalton's atomic theory was based on the 
following ideas: 

1) That all matter consists of extremely small particles, called 

2) That all the atoms of any one element are alike in size, shape, 
and weight. 

3) That the atoms of one element differ from those of all other 
elements in size, shape, and weight. 

4) That chemical changes are changes in the combination of atoms 
with each other. 

5) That even in the most violent chemical change, the indivisible 
atoms do not break into pieces. 

The diameter of the hydrogen atom is about 1/250,000,000 inches. 
This is several hundred times smaller than the average-sized bac- 

To explain his theory, Dalton drew pictures of atoms. Each atom 
was represented by a circle. Since the atoms of elements are unlike, 
he varied the pictures of the circles as follows: 

Hydrogen Oxygen Nitrogen 

Carbon Sulfur Q Gold 

Dalton pictured chemical change as the union of one or more 
atoms of one element with atoms of other elements. He believed that 
when mercury is heated in air, one atom of mercury unites with one 
atom of oxygen, forming a particle of the compound, mercuric oxide. 
To demonstrate this union of atoms, Dalton constructed model 
spheres, bringing them into contact with each other: 

1 atom of liquid 1 atom of gaseous 1 molecule of red 

mercury "*" oxygen * mercuric oxide 

According to Dalton, atoms preserve their individuality in all 
chemical changes. Hence, Dalton described an atom as the smallest 
part of an element that takes part in a chemical change without itself 
being altered. Atoms combine to form molecules. Two or more atoms 
may combine into a molecule of an element or into a molecule of a 
compound. A molecule is the smallest part of a compound or element 
that has the chemical properties of that compound or element. 


Inertia in scientific thinking. Dalton's theory was strongly at- 
tacked by the leading scientists of his day. One of the most eminent 
of them said he could not understand "how any man of sense or sci- 
ence would be taken in by such a tissue of absurdities." Dalton's 
theory was the result of creative imagination and the boldness of a 
great thinker. Dalton had never seen nor weighed an atom. Yet his 
theory was of practical value and was accepted gradually by the scien- 
tific world as a useful working hypothesis by which chemical changes 
could be explained. 

The newer electron theory of matter (see Chapter 11) has ex- 
panded Dalton's theory. It has been modified in details, but its gen- 
eral applications still hold. 

The use of Dalton's theory. One of the questions that scientists of 
Dalton's time were studying was whether the composition of a com- 
pound is always the same or whether it varies. Some believed that 
compounds are always formed from fixed amounts of elements. They 
believed therefore, that the composition of a compound is always the 

The French chemist, Claude Berthollet (ber'to'le') ran some ex- 
periments to test this question. On the basis of his experiments he be- 
lieved that the composition of compounds might vary to some extent. 

Joseph Proust (prdost) , another Frenchman, set out to settle this 
difference of opinion. He repeated Berthollet's experiments, using 
the purest of chemicals and the most delicate apparatus available. 
Taking every precaution to prevent error, he found mistakes in 
Berthollet's work. He found that his fellow-scientist had used im- 
pure compounds and substances such as glass and mixtures of metals 
(alloys) and mixtures of liquids, which were not pure compounds. 

For eight long years, the difference of opinion persisted. Never, 
however, did it become anything but an honest, truth-seeking discus- 
sion. Personal whims and prejudices did not decide the matter. 
When Berthollet considered Proust's evidence and discovered his 
own errors, he accepted Proust's verdict and agreed that the compo- 
sition of compounds is always the same. 

The law of definite proportions. This law states that the elements 
in a compound always occur in a definite proportion by weight. This 
is another way of saying that the composition of compounds is always 
the same. 

Dalton's little circles very neatly explained the law. The weight 
of atoms of any element is always the same. Compounds are com- 
posed of these minute and unchangeable atoms. Therefore the com- 
position of compounds by weight must be definite and uniform. 



Dalton discovers the law of multiple proportions. Dalton knew 
that one atom of carbon unites with one atom of oxygen O to 
produce the deadly gas, carbon monoxide O In this compound 
the carbon weighs f as much as the oxygen. This fraction can be ex- 
pressed as the ratio of 3 to 4. 

Carbon also combines with oxygen to form carbon dioxide. Dal- 
ton wrote this combination as 0QO- In this compound the carbon 
weighs only f as much as the oxygen. The ratio is three parts of car- 
bon to eight parts of oxygen: 3 to 8. From observing this and other 
similar combinations, Dalton formulated another fundamental law 
of chemistry, the law of multiple proportions. 

You note that in both carbon monoxide, CO, and in carbon di- 
oxide, CO 2 , the weight of the carbon is the same. But the weight of 
the oxygen in carbon monoxide is 4, and in carbon dioxide is 8. 
Thus three parts of carbon combine with either 4 parts or 8 parts of 

When two elements combine to form more than one compound, 
with the weight of one element remaining fixed, the ratios of the 
weights of the other elements are small whole numbers. 

Thus the amounts of oxygen that unite with three parts of carbon 
are in the ratio of 4 to 8, or 1 to 2. 

How the discovery of hydrogen peroxide helped to uphold the law 
of multiple proportions. In 1818, Louis Thenard (ta nar') , a French 
teacher of chemistry, discovered a compound, which upon analysis 
was shown to be made up of equal volumes of oxygen and hydrogen. 
This compound is hydrogen peroxide. 

Hydrogen and oxygen combine to form two different compounds. 
Water is composed of one part hydrogen and eight parts oxygen by 
weight. Hydrogen peroxide is composed of one part hydrogen and 
sixteen parts oxygen. Thus the ratio of the weights of oxygen that 
combine with a fixed weight of hydrogen is 8 to 16, or 1 to 2. 





( Water H 2 O 
(^ Hydrogen peroxide H 2 O 2 

1 H 
1 H 

80 "I 

160 / 

1 to 2 

f Carbon monoxide CO 
^ Carbon dioxide CO 2 


40 ) 

8O ) 

1 to 2 

This maze of tanks and pipes is re- 
quired for the commercial preparation 
of hydrogen peroxide. 

Properties and uses of hydrogen peroxide, FLO,. Since water and 
hydrogen peroxide have different chemical compositions, they have 
different physical and chemical properties. Hydrogen peroxide is a 
colorless liquid, about one and one-half times as heavy as water. It 
is odorless and mixes with water, alcohol, or ether. It is useful com- 
mercially because it is unstable. That is, heat or light decomposes it 
easily into water and oxygen: 

H 2 O 2 -* H 2 + O T 

The oxygen atom that is liberated is in a very active state, i^eady to 
combine with another atom of oxygen or with any other substance 
at the instant of liberation. This very active atomic oxygen is some- 
times called nascent (newborn) . Nascent oxygen is written as (X 
ordinary oxygen as O 2 . The arrow pointing upwards represents a gas. 

Some colored compounds lose their color when they are oxidized. 
Fibers containing compounds which give them their color can be 
bleached by exposing them to nascent oxygen. Hydrogen peroxide 
is used as an oxidizing agent to bleach, or decolorize, cotton goods, 
wool, wood pulp, wood used for furniture, as well as silk, hair, feath- 
ers, glue, and other animal substances. 

Some bacteria are destroyed when exposed to oxygen. Hydrogen 
peroxide is therefore used as a household antiseptic. The household 
product actually is mainly water with a small amount (usually 3 per 
cent) of hydrogen peroxide dissolved in it. It also contains some sub- 
stance such as acetanilid, which retards the decomposition of the hy- 
drogen peroxide. As an antiseptic, hydrogen peroxide is not very 
effective since the oxygen it releases does not reach enough of the 




Commercial preparation of hydrogen peroxide. It barium perox- 
ide, a white solid, is treated with dilute sulfuric acid at a temperature 
below L r )C., hydrogen peroxide and barium sultate are formed. Bar- 
ium sulfate is a white insoluble solid which settles out. 

Sulfuric acid -f barium peroxide > hydrogen peroxide 4- barium sulfate 
;H 2 !SO4~~~+"~~~BaO 2 j -> H 2 O 2 + BaSO 4 1 

This type of reaction is called double replacement barium replaces 
hydrogen. The arrow pointing downward after BaSO 4 indicates that 
this compound is insoluble and separates out, or precipitates. The 
insoluble substance that separates out is called a precipitate. 

Most of the hydrogen peroxide produced commercially is made by 
gently heating pcrsulfuric acid, H_,S,O N , which reacts with water: 

Persulfuric acid 



2H 2 

sulfuric acid 
2H 2 S0 4 

+ hydrogen peroxide 
+ H 2 O 2 

In this process the hydrogen peroxide is distilled out. Superoxol is a 
30% H 2 O 2 solution. A 90% solution of this substance is used as a 
rocket fuel. 

The Glenn Martin Company 

Off to an altitude of 160 miles roars 
the Navy's Viking No. 1 1 rocket. 
In flight this rocket will attain a 
speed of 4300 miles per hour. 
Rocket development depends 
greatly upon fuel research carried 
on in chemical laboratories. 



Harrow, Benjamin. The Romance of the Atom, pp. 27-33. 
Boni & Liveright, New York, 1927. Origin and development 
of ideas about atoms. 

Langdon-Davies, John. Inside the Atom. Harper & Bros., 
New York, 1933. Amusing, popular introduction to science 
and the nature of matter. 

Leicester, H. M. and Klickstein, H. S. Source Book in Chem- 
istry, pp. 215-220. McGraw-Hill Book Company, New York, 
1952. Gives John Dal ton's observations on the constitution of 

Thomson, J. Arthur. The Outline of Science, pp. 245-253. 
G. P. Putnam's Sons, New York, 1937. "Foundations of the 
Universe." In "The World of Atoms" the size and energy of 
these tiny particles are very simply discussed. 


1. The idea of individual indivisible particles of matter 
originated with the ancient Greek teachers. Dalton used this 

[idea and developed the atomic theory in 1803. 

2. The chief assumptions of Dalton's theory were: a) That 
all matter consists of extremely small particles, called atoms, 
b) That all the atoms of any one element are alike in size, 
shape, and weight, c) That the atoms of one element differ 
from those of all other elements in size, shape, and weight, 
d) That chemical changes are changes in the combination of 
atoms with each other, e) That even in the most violent 
chemical changes, the indivisible atoms do not break into 

3. An atom is the smallest part of an element that takes 
part in a chemical change without itself being altered. 

4. A molecule is composed of two or more atoms and is the 
smallest part of a compound or element that has the proper- 
ties of that compound or element. 

5. Personal whims, prejudices, or prestige should play no 
part in settling scientific problems. 

6. The law of definite proportions states that the elements 
in a compound always occur in a definite proportion by 

7. The law of multiple proportions states: When any two 
elements combine to form more than one compound, with the 
weight of one element remaining fixed, the ratios of the weights 
of the other element are small whole numbers. 




Group A 

1. When did John Dalton advance his atomic theory? 

2. Was the idea of atoms new in Dalton's time? Explain. 

3. How did Dalton's approach to the study of the nature of 
matter differ from that of the Greeks? 

4. State the five essentials of Dalton's atomic theory. 

5. How did Dalton distinguish between atoms of different 

6. According to Dalton, what happens when elements 

7. How did the scientific world receive Dalton's theory? 

8. Describe the difference of opinion between Berthollet 
and Proust. 

9. (a) State the fundamental law of chemistry that Dalton 

(b) Show how it applies to the composition of CO and 
C0 . 

10. (a) When and (b) by whom was H 2 O 2 prepared for the 
first time? 

1 1 . Compare the physical properties of H 2 O and H 2 O 2 . 

12. Contrast the chemical properties of H 2 O and H 2 O 2 . 

13. Compare the compositions of H 2 O and H 2 O 2 by volume 
and weight. 

14. State two ways in which the decomposition of H 2 O 2 may 
be retarded. 

15. Write the word-equation for the decomposition of H 2 O 2 . 

16. What is the difference between O 2 and nascent oxygen? 

t . . . 

17. Describe the two most useful properties of H 2 O 2 . 

18. What is the great advantage of H 2 O 2 as a bleaching 

19. What is an antiseptic? 

20. State a commercial method of preparing H 2 O 2 . 

21. Write the equation for the preparation of H 2 O 2 from 
BaO . 


22. What is a precipitate? 

23. Compare the chemical properties of O 2 , O 3 , and O. 

Group B 

24. Explain how the law of multiple proportions is based 
on Dalton's atomic theory. 

25. Two oxides of nitrogen, nitrous oxide (N.,O) and ni- 
trogen dioxide (NO 2 ) , show the following ratios by weight: 
In nitrous oxide, the ratio by weight of nitrogen to oxygen 
is 7 to 4; in nitrogen dioxide, the ratio by weight of nitrogen 
to oxygen is 14 to 32. With what fundamental law of chemistry 
are these figures in accord? 

26. Do we still believe that all atoms of the same element 
weigh the same? 

27. (a) Is Dalton's atomic theory still a theory or has it been 
proved experimentally? (b) Explain your answer. 

28. Dalton's hobby of recording weather conditions was 
greatly responsible for the atomic theory. Can you cite another 
example in science where a hobby has resulted in a great 

29. Write a 2- or 3-page report on the life of John Dalton. 


1. Dalton's atomic theory is a beautiful example of creative 
imagination in pure science. Until very recently America, al- 
though it excelled Europe in inventions and applied science, 
has lagged behind the Old World in the kind of creative imagi- 
nation represented by Dalton's theory and the germ theory of 
disease which Pasteur gave humanity. Can you give reasons 
for this state of affairs? What of the future? 

2. Add a pinch of manganese dioxide to one-third of a test 
tube of hydrogen peroxide from your medicine cabinet. Insert 
a glowing splint. Prepare a brief report on the reaction which 
takes place. 

3. Prepare a debate or a class discussion on the topic, "If 
prejudices and superstitions did not exist, social change would 
proceed more rapidly." 

4. Form a committee to investigate and report on prejudices 
and superstitions common to your community. Do any of these 
have a scientific foundation? 

5. Make models of the atoms of hydrogen, carbon, sulfur, 
gold, mercury, and nitrogen as Dalton constructed them, using 
marbles or some other suitable spheres. 




The mysterious symbols of alchemy. Ancient alchemists feared 
that they would lose their unique social position if anyone could 
understand their work. Consequently they used strange symbols, 
both to conceal the true nature of their writings and also to give 
themselves an added air of mystery and magic. They represented 

sand by .jjj.^ , glass by Q Q, soap by ^^-, and salt by . To 
^^^^^ ^"^ 

them the symbcJ for perfection and also for the sun was the circle 
O Hence it was used to represent sun-colored gold, which they con- 
sidered the perfect metal. 

Many of their symbols were derived from ancient mythology. The 
lance and shield of Mars, god of war, f , represented, appropriately 

enough, iron. The looking glass of Venus, O, was their symbol for 
copper because, according to legend, Venus had first appeared on the 
shores of the island of Cyprus, long famous for its copper mines. 

Often an alchemist would develop a set of symbols for his own 
special use and would reveal the meaning to no one, not even to his 
brother alchemists. In one Italian manuscript written in the seven- 
teenth century, the element mercury was represented by 20 different 
symbols and 35 different names. As long as alchemy was a purely 
personal practice carried on for selfish ends, the confusing symbolism 
presented no problem. However, when the universal science of 
modern chemistry began to emerge, it was essential to develop a 
system of symbols which could be understood easily in every country 
of the world. Dalton, and a few others before him, had attempted 
to substitute some reasonable system for the jungle of weird signs 
and strange names used by the alchemists. However, these attempts 
failed largely because they were unwieldy and inconvenient. 



Berzelius helps bring order out of chaos. Jons Berzelius (ber- 
ze'li-ws) was a Swedish orphan who became, in the words of Sir 
Humphry Davy, "one of the great ornaments of his age." Berze- 
lius invented a simple system of chemical notation which he intro- 
duced in 1814. Today it is used in every country of the world. 

Berzelius said very logically: "It is easier to write an abbreviated 
word than to draw a figure. The chemical signs ought to be letters 
for the greater ease of writing and not to disfigure a printed book. 
I shall therefore take for the chemical sign the initial letter of the 
Latin name of each chemical element, thus, C, H, N, O, S, and P. 
If the first letter be common to two metals, I shall use both the initial 
letter and another letter they have not in common, as gold (aurum) , 
Au; silver (argentum) , Ag; antimony (stibium) , Sb; tin (stannum) , 
Sn." The first letter of the symbol is always capitalized; the second, 
if there is a second, is not. 

How a compound is represented by a formula. When we are deal- 
ing with a compound rather than an element, we write its abbrevi- 
ated form, or formula, by placing side by side the symbols of the 
elements that compose the compound. For example, the formula of 
zinc oxide, a compound of zinc and oxygen, is ZnO; the formula of 
hydrochloric acid, a compound of hydrogen and chlorine, is HC1. 
A formula not only represents the name of a compound, but also 
one molecule of that compound. Similarly a symbol represents both 
the name of an element, and one atom of that element. 

The use of the subscript. Berzelius used a subscript, a small num- 
ber placed below and to the right of a symbol, to indicate the number 
of atoms of the element represented by that symbol. Thus, H 2 O 
represents one molecule of water containing two atoms of hydrogen 
and one atom of oxygen. The subscript 1 is never written. 

As you use Berzelius' system, you will see its remarkable simplicity 
and great value. However, it was not accepted without a struggle. 
Even Dalton protested against it saying, "Berzelius' symbols are 
horrifying. A young student might as soon learn Hebrew as make 
himself acquainted with them." Evidently, Dalton must have for- 
gotten his own very complex pictures of various compounds. 

How is the proper subscript determined? You may have already 
wondered how it is possible to tell what subscript to use. Why A1 2 O 8 
and ZnCl 2 and not A1O and ZnCl? Must we memorize all formulas or 
are there definite rules to guide us? 

Fortunately, it is possible to write formulas without first mem- 
orizing them. To do so, however, we must have a thorough under- 
standing of valence. 



What is valence? The law of definite proportions tells us that 
atoms occur in compounds in fixed ratios. The following table shows 
the ratio in which hydrogen combines with other elements to form 
four common compounds. 




HCI (hydrogen chloride) 
H 2 O (water) 
NH 3 (ammonia) 
CH* (methane) 



1 chlorine atom 
1 oxygen atom 
1 nitrogen atom 
1 carbon atom 

In this table, note the difference in the number of hydrogen atoms 
with which one atom of the other elements combines. The valence 
of an element is the number that tells us how many atoms of hydro- 
gen normally combine with one atom of that element. This is a 
simplified definition, but will serve us for the present. 

From the table we see that one atom of chlorine combines with 
one atom of hydrogen to form HCI. Thus, we say that chlorine has a 
valence of one, or as chemists put it, is monovalent. One atom of 
oxygen combines with two atoms of hydrogen. So we say that oxygen 
has a valence of two, or is divalent. Nitrogen has a valence of three, 
or is trivalent. Carbon has a valence of four, or is tetravalent. 

The idea of valence was introduced in 1852 by Edward Frank- 
land, an English chemist. Hydrogen is used as the standard, because 
its atom never combines with more than one atom of any other ele- 
ment. Hence, if its valence is considered to be one, the valence of 
every other element must be a whole number. 

Learning valences. There is no "royal road" to the study of va- 
lence. You will find that memorizing the valences of the more com- 
mon elements will save you a tremendous amount of work and will 
give you a better understanding of the material to follow. Table 2 
lists the valences which are considered essential. 

Notice th'at in the table, the symbols of the metals are followed 
by plus (+) signs and the nonmetals and radicals followed by 
minus signs ( ) . These signs represent electric charges, for reasons 
which you will learn in Chapter 16. The number of these electric 
charges correspond to the valences of the element or radical. Group- 



ing elements as metals and nonmetals is not a completely satisfactory 
means of classification because some elements behave at times as 
metals and at other times as nonmetals. 

Binary compounds. A binary compound is composed either of two 
elements, two radicals, or one element and one radical. If you keep 
in mind the following rules, you will find writing the formulas for 
binary compounds is a simple process. Practice is essential. 










*Ammonium NH 4 * 
Copper (cuprous) Cu + 
Lithium Li * 
Mercury (ous) Hg * 
Potassium K 
Silver Ag + 
Sodium Na + 

Barium Ba "*"*" 
Calcium Ca++ 
Copper (cupric) Cu ** 
Iron (ferrous) Fe ** 
Magnesium Mg" 1 " 1 " 
Mercury (ic) Hg*^ 
Zinc In** 

Aluminum Al *+ + 
Antimony Sb" 1 "*"* 1 
Arsenic As * <f+ 
Chromium Cr "*"*"*' 
Iron (ferric) Fe "*"*"*" 


Bromine (bromide) Br~ 
Chlorine (chloride) Cl~ 
Fluorine (fluoride) F " 
Iodine (iodide) l~ 

Oxygen (oxide) O 
Sulfur (sulfide) S~" 

Nitrogen N 
Phosphorus P 


*Acetate C 2 H 3 O 2 ~ 
Bicarbonate HCOs" 
Chlorate CI0 3 " 
Hydroxide OH~ 
Nitrate NOa " 
Nitrite NO 2 - 

Carbonate COs 
Sulfate S0 4 ~~ 
Sulfite SO 3 

Phosphate PO 4 

* Radicals: A radical is a group of atoms acting as a single atom and having its awn individual 
valence. The ammonium radical, NH^ has a valence of ?. 



1 ) Write the symbol with a positive valence first, followed by 
the symbol with a negative valence. Add plus and minus signs 
to the upper* right of the symbols, that is, in superscripts, to 
show the valence for each symbol. 

2) // the valences of the symbols are equal no subscripts are 
added. This rule is followed unless the subscripts represent 
the actual structure of the compound. Thus, the formula of 
hydrogen peroxide is H 2 O 2 and not HO, since a molecule of 
hydrogen peroxide actually contains two atoms of hydrogen 
and two atoms of oxygen. 


3) Since every compound is electrically neutral, the number 
of positive charges must be the same as the number of its 
negative charges. Therefore, // the valences are not equal in 
numerical value, subscripts must be added to equalize them. 
Add to each symbol a subscript of the same numerical value 
as the valence of the other symbol. The subscript 1 is never 

4) A radical acts like an element, that is, it usually passes 
thiough a chemical reaction unchanged. It should be placed 
in parentheses only if it is followed by a subscript greater than 

EXAMPLE A: Write the formula for the compound, zinc oxide. 

1) Zinc is written Zn and has a valence ot plus two; oxygen 
is written O and has a valence of minus two. The symbol lor 
zinc appears first in the formula since zinc is a metal and oxy- 
gen is a nonmetal. Indicate valences by plus and minus signs. 


2) Since the valences are equal no subscripts are written, 
and the subscript for each symbol is understood to be one. 

3) The proper formula for zinc oxide therefore is ZnO. 

EXAMPLE B: Write the formula for the compound, cupric 

1) Copper is written Cu and has a positive valence of two; 
chlorine is written Cl and has a negative valence of one. Write 
the symbol for copper first because it has the positive valence. 
Indicate the valence of each element by using plus or minus 
signs. (Note that the cupric valence of copper should be used. 
See Table 2.) 


2) The subscript of each symbol must be equal to the va- 
lence of the other symbol. Since Cu has a valence of plus two, 
give Cl the subscript two. Since Cl has a valence of minus one, 
the subscript for Cu is understood to be one. The crossed ar- 
rows show these relationships. 

The negative and positive valences are equal because there is 
one atom of copper with a valence of plus two, and two atoms 
of chlorine each with a valence of minus one. 

3) The proper formula for cupric chloride is therefore 


EXAMPLE C: Write the formula for the compound, magnesium 

1 ) Magnesium (Mg) has a valence of plus two. The sulfate 
radical (SO 4 ) has a valence of minus two. Therefore, Mg ap- 
pears first in the formula. Indicate the valences by using plus 
and minus signs. 

Mg++SO 4 

2) The valences of the two symbols are equal. Therefore, 
the subscript of each is understood to be one, and no sub- 
scripts are written. 

3) The proper formula for magnesium sulfate is MgSO 4 . 

EXAMPLE D: Write the formula for the compound, zinc nitrate. 
I ) Zinc (Zn) has a valence of plus two. The nitrate radical 
(NO 3 ) has a valence of minus one. Therefore, Zn appears first 
in the formula. Indicate the valences by using plus and minus 

Zn++NO 3 - 

2) Since Zn has a valence of plus two, we give NO 3 the sub- 
script two and enclose it in parentheses. NO 3 has a valence of 
minus one and we consider Zn to have the subscript one. The 
crossed arrows show these relationships. 

3) The proper formula for zinc nitrate is Zn(NO 3 ) 2 . 

Some elements have more than one valence. Iron is divalent 
in ferrous compounds and trivalent in ferric compounds. Fer- 
rous chloride provides an example of divalent iron: Fe++Cl 2 -. 
Ferric chloride provides an example of trivalent iron: 
Fe+++Cl 3 -. 

Mercury is monovalent in mercurous compounds and di- 
valent in mercuric compounds; copper is monovalent in 
cuprous compounds and divalent in cupric compounds. What 
is the meaning of -ous and -ic as related to valence? 

In the formula Fe 3 O 4 (Magnetic oxide of iron) you might 
think that iron has a valence of four and oxygen a valence of 
three. However, the real explanation is that this compound 
is a combination of Fe++O~-, in which Fe is divalent, and 
Fe a +++O 8 - ~, in which Fe is trivalent. Oxygen is always divalent. 
Or, Fe 3 6 4 may be thought of as FeO Fe 2 O 3 . From these 
formulas it is apparent that some elements may exhibit two or 
more valences. The electron theory offers an interesting 


explanation of the fact that certain elements have more than 
one valence (see Chapter 11). 


1. Write the formulas for the following compounds showing 
the + an d signs and the arrows pointing toward the sub- 
scripts: (a) sodium chloride, (b) calcium bromide, (c) ferric 
iodide, (d) potassium fluoride, (e) barium oxide. 

2. Write the formulas of: (a) magnesium chloride, (b) zinc 
oxide, (c) aluminum nitride, (d) potassium sulfide, (e) cu- 
prous chloride. 

3. Write the formulas of: (a) aluminum chloride, (b) arse- 
nic oxide, (c) calcium phosphide. 

4. Write the formulas of: (a) sodium hydroxide, (b) potas- 
sium sulfate, (c) mercurous phosphate. 

5. Write the formulas of: (a) calcium bicarbonate, (b) cu- 
prous carbonate, (c) magnesium phosphate. 

6. Write the formulas of: (a) aluminum nitrate, (b) ferric 
sulfate, (c) chromium phosphate. 

7. Carbon (C++++) and silicon (Si++++) are tetravalent. 
Write the formulas of: (a) carbon tetrachloride and (b) sili- 
con dioxide. 

8. Write the formulas of: (a) mercuric nitrate, (b) sodium 
nitrate, (c) mercuric chlorate, (d) mercuric hydroxide, (e) 
mercuric carbonate, (f) mercurous sulfate, (g) calcium sul- 
fite, (h) mercuric phosphate, (i) mercurous chloride. 

How to determine valence in compounds of more than two ele- 
ments. Remembering that in every compound the number of posi- 
tive charges must equal the number of negative charges, let us find 
the valence of Cr in K 2 CrO 4 . There are four oxygen atoms each with 
two negative charges making a total of eight negative charges. To bal- 
ance these, the compound has two potassium atoms each having one 
positive charge or a total of two positive charges. The compound 
must therefore have six more positive charges which must come from 
the metal chromium. Hence the valence of Cr in this compound is 
plus six. 

Chemistry has a language and nomenclature of its own. Lavoisier 
realized the importance of language to a science. In 1789, the year 
in which the Bastille was stormed, he published a book written in 
the new language of chemistry. It did not contain the obscure words, 
the mystic symbols, and the pompous phrases of alchemy. 

An outstanding nuclear scientist is Dr. Glenn T. 
Seaborg of the Radiation Laboratory, University of 
California, Berkeley, California. He has played an 
important role in the discovery of the trans-uranium 
elements, numbers 93-100. What names have 
been given elements 93-98? Can you suggest 
their derivation? 

University <>! Call In 

In naming the elements, several methods were used. Some, in- 
cluding bromine, meaning stench, were named after a physical prop- 
erty. Some, including argon, meaning idle, were named after a chem- 
ical property. Some, including /;o/onium, germanium, gallium, and 
americium, were named after countries or other geographic regions. 
Some were named after the city or state connected with their dis- 
covery. Thus hafnium was christened after the Latin name for the 
city of Copenhagen, where it was discovered. Radioactive curium was 
named after the Curies, who were the earliest investigators in the field 
of radioactivity. 770rium and tantalum were named after figures 
in mythology. The origins of the names of the earliest known ele- 
ments have been lost in the darkness of antiquity. 

Names of metals and metallic radicals usually end in either -turn 
or -um, as sodium, potassium, platinum, curium, hafnium, alumi- 
num, calcium, and the ammonium radical. 

Names of nonmetals and nonmetallic radicals usually end in -ine 
or -gen, as chlorine, bromine, iodine, oxygen, and nitrogen. 

The most common suffixes used in naming compounds are -ide, 
-ate, -ite, -ous, and -ic. The suflix -ide represents a binary compound. 

The suffixes -ate and -ite indicate, as a rule, compounds of three 
elements, one of which is oxygen. The -ite compound contains fewer 
oxygen atoms than the corresponding -ate compound. Thus sodium 
sulfate, Na.,SO 4 , contains four atoms of oxygen, and sodium sulfite, 
Na.,SO 3 , three atoms of oxygen. The -ate compounds are salts (see 
page 197) of -ic acids, and -ite compounds are salts of -ous acids. 

The suffixes -ous and -ic indicate compounds in which the metal 
has a lower valence in the case of -ous and a higher valence in the case 
of -ic. An -ous acid contains fewer oxygen atoms than an -ic acid; 
thus sulfurous acid is H 2 SO 3 and sulfuric acid is H 2 SO 4 ; chlorous 
acid is HC1O., and chloric acid is HC1O 3 . 



Some commonly used prefixes in chemistry are mono- (or uni-) , 
di- (or hi-) , tri-, tetra-, and pent-. Mono- (or uni-) , di- (or hi-) , tri-, 
tetra-, and pent- stand for one, two, three, four, and five atoms. Thus, 
CO is carbon monoxide, and CO 2 is carbon dioxide. P 2 O 3 is phos- 
phorus trioxide and P 2 O 5 is phosphorus pentoxide. The prefix per- 
means more atoms of an element than are found in a more common 
compound, and the prefix hypo- means less atoms. Thus chloric acid 
is HC1O 3 , perchloric acid is HC1O 4 ; chlorous acid is HC1O 2 , and 
hypochlorous acid is HC1O. 

In organic chemistry, a branch of chemistry which deals with the 
more than 650,000 compounds of carbon, a more comprehensive 
nomenclature has been carefully worked out. We shall learn more 
about organic chemistry later. 


Caven, Robert M. and Cranston, John A. Symbols and 
Formulae in Chemistry, pp. 1-29. Blackie &: Son, London, 
1928. Development and use of symbols and formulas are traced 
with great clearness and interest in this valuable work. 

Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 136- 
156. Simon and Schuster, New Ybrk, 1948. "A Swede Tears Up 
a Picture Book" deals with Berzelius' life and his contributions 
to the development of chemistry. 

Kendall, James. At Home among the Atoms. D. Appleton- 
Century Co., New York, 1932. This eminent chemist intro- 
duces in novel form the problem of valences. He calls the 
chapter "Valencia." 

Oesper, Ralph E. "The Birth of Modern Chemical Nomen- 
clature." Journal of Chemical Education, June, 1945, pp. 290- 
292. Aii interesting story. 


1. Each element has a symbol that stands for one atom of 
the element. 

2. Each chemical compound may be written in abbreviated 
form as a formula by placing the symbols of the elements that 
compose the compound side by side. 

3. A radical is a group of elements that act as a single ele- 

4. Subscripts ate used in a chemical formula to indicate the 
number of atoms of any element which occur in a molecule of 


any substance. Subscripts are also used following the paren- 
theses around radical groups to indicate how many radicals 
are present in one molecule. 

5. The valence of an element is a number that represents 
the number of atoms of hydrogen with which one atom of that 
element normally combines in forming a compound. 

6. Names of metals and metallic radicals usually end in -turn 
or -urn; names of nonmetals often end in -ine or -gen. 


Group A 

1. What is a symbol? 

2. Who introduced the modern symbols and formulas of 

3. Why was it necessary to replace the old alchemical 
symbols with a new system of symbols? 

4. Write the symbols of three gaseous, one liquid, and 
three solid elements. 

5. Several elements have the same initial letter. How do 
we indicate these elements by symbols? 

6. What does a formula indicate? 

7. (a) Define valence, (b) Who introduced valence into 

8. What element having a valence of one is used in de- 
termining the valence of other elements? 

9. What kind of element would have a valence of zero? 
10. Of what importance to you is a knowledge of valence? 

1 1. What is a radical? 

12. Copy and complete the following statements: Metals 
have . . . valences; nonmetals have . . . valences. The most 
common radical with a plus valence is .... 

13. Copy and complete the following statements: Three 
metals with a valence of one are . . . , . . . , and .... Three 
divalent metals are . . ., . . ., and .... Three monovalent non- 
metals are . . . , . . . , and .... The bicarbonate radical has a 
valence of .... A radical with a valence of three is .... 

14. State four rules for writing formulas. Illustrate each. 



15. Give two examples of elements that have more than 
one valence. 

16. Copy and complete the following table. Do not write 
in this book. 

Bromide Sulfide Chlorate Sulfate Phosphate Oxide Hydroxide 







17. Make a list of the formulas of the phosphates of eight 
different metals. 

18. Make a list of the formulas of the carbonates of eight 
different metals. 

19. Make a list of the formulas of five ammonium com- 

20. Correct the following: FeCl, CuS 2 , Ag(NO 3 ) 2 , KSO 4 , 
Na 2 C10 3 , and NH 4 (OH) . 

21. Write the names and formulas of five compounds of 

It ... 

22. Name five elements and tell how their names were de- 

23. How can you tell whether the element ruthenium is a 
metal or a nonmetal? 

24. Give the formulas and names of two acids that illustrate 
the difference in the use of the suffixes ~ous and -ic. 

25. Give the formulas and names of two compounds other 
than acids that illustrate the difference in the use of -ous 
and -ic. 

I . . 

26. How would you name the two compounds, BaO and 

Ba0 2 ? 

27. Explain the meaning of each letter and subscript in 
these formulas: HNO 8 , FePO 4 , Na 2 CO 8 . 

28. Mark the valences of the elements and radicals in these 
compounds, using -{- and signs. CuSO 4 , HgCl 2 , NaClO 3 , 
Ca (HC0 3 ) 2 , X 3 (P0 4 ) 2 , (NH 4 ) 8 Y, MnO 2 . 


29. Determine the valence of: (a) sulfur in H 2 SO 4 , (b) 
manganese in KMnO 4 , and (c) chromium in K 2 Cr 2 O 7 . 

Group B 

30. An unknown element X has a valence of three. Write 
the formula for its oxide. 

31. The metal Y has a valence of two and the nonmetal Z 
has a valence of two. Write the formula for their compound. 

32. The nomenclature of chemistry is still not completely 
organized. Can you give any reasons for this state of affairs? 

33. What element was named after a Finnish chemist? 
(Read the chapter on the rare-earth elements in Weeks' Dis- 
covery of the Elements.) 

34. Write a two- or three-page report on the life of Berzelius. 
See list of additional reading material. 


1. Make a chart and, with the help of your teacher of art, 
include as many alchemical symbols as you can and their 
modern equivalents. 

2. The old alchemists wanted to keep their knowledge hid- 
den from the rest of the world, and so used strange symbols 
and mysterious language, (a) Name two groups of people who 
are modern equivalents of the alchemists, (b) Tell how they 
keep their knowledge to themselves and clothe their activities 
with a veil of mystery, (c) What are the reasons for their 

3. Form a committee to make a report on the trades, pro- 
fessions, and businesses in your community that have "lan- 
guages of their own." (a) What reasons have they for these 
languages? (b) What, after all, is the function of language? 



. . . It ought to be esteemed much 
less disgraceful to quit an error for 
a truth than to be guilty of the van- 
ity and perverseness of believing a 
thing still, because we once believed 
it. Robert Boyle, 1627-1691 

Exploring the atmosphere. Every day many of us see gleaming air- 
planes flying swiftly overhead. What are conditions up in the air 
where they travel? Of what is the air in which they travel composed? 

Have you ever seen the glowing trail of a meteor as it flashed 
across the sky? Why did it burn? Have you ever examined the fused, 
fire-scarred surface of a meteorite? Was it hot when it entered our 
atmosphere, or did it become heated as it traveled through it? 

Probably you know the answers to these questions, but primitive 
men would not have been able to answer them. The most learned of 
the alchemists, even as late as the seventeenth century, did not have 
the answers. Some of them believed that air was empty space without 
weight and without substance, while others believed air to be one of 
the "four elements." 

It was not until 1643 that Torricelli (tor-re-chel'le) invented the 
mercury barometer, which shows that air has weight and exerts pres- 
sure. We know today that approximately 15 pounds (14.7 Ib.) of 
air rest upon every square inch of the surface of the earth. Upon 
every square foot rests a column of air weighing more than one ton! 

The density of the atmosphere is not the same throughout its 
entire extent. It is densest at the surface of the earth, and the farther 


the surface of the earth, the lighter the air becomes, At five 
above the earth, air pressure Is only 5.5 pounds. At this eleva- 
tion the engine of an airplane cannot get sufficient oxygen to burn its 
fuel. U is half-starved. What it needs to function properly is more 
oxygen, The turbosupercharger supplies this by foreleg more air 
into the It has been calculated that even at a height of 2000 

miles there 'is still some air, although above five miles normal breath- 
ing of human is impossible. About 90 percent of the weight 
ofthf is within 12 miles of the earth, 

With respect to temperature, the atmosphere Is divided into sev- 
eral layers, 'The lowest layer, about six miles deep, is called the 
troposphere. Within the troposphere the temperature falls about 
: 1F. with each, rise of 300 feel. The layer above this is called the. 
! stratosphere where temperature from about 30 to 65F. 

I In the next layer, which lies between 18 and 28 miles above the 
earth's surface, a temperature of about 65 degrees below zero oi^the 
Fahrenheit scale is almost constant. Above these layers, at a height 
of between 50 and 200 miles, are several layers of electrified particles 
(of which the Kennelly-Hcaviside and the Appleton were the earliest 
discovered) , which prevent most radio waves from out into 

space, and" reflect them back to the earth on their way around the 
globe. This is called the ionosphere, Its from. 

A of a cross-section of the showing fhit approximate relative" 

osition of the various layers. 


morning to night and varies with the seasons. The temperature in 
these ionized bands rises with elevation. 

Extremely short radio waves of sufficient energy can penetrate 
these electrified layers, as was proved in 1946 by Army Signal Corps 
scientists, who succeeded in sending radar impulses to the moon. 
Reflections from the moon were received. The round trip, 477,714 
miles, was made in 2.4 seconds. 

What is air? Air is made up of a number of different gases. During 
the eighteenth century three of these nitrogen, oxygen, and car- 
bon dioxide were isolated, or obtained separately, from it. Indi- 
cations were noted that still other gases, less easily obtained in a 
pure state, might possibly be present in small quantities. These 
gases were actually found later on. 

Is air a mixture or a compound? How are the gases of the air held 
together? Are they chemically united or are they merely mixed, just 
as sand might be mixed with clay? To answer these questions, sci- 
entists made use of the law of definite proportions. If air is a com- 
pound, they reasoned, then the composition of air must be constant. 

To determine whether or not the composition of air is constant, 
Dalton and others collected and analyzed samples of air taken from 
thousands of different places from the tops of mountains, over 
lakes, in valleys, in sparsely settled regions, and in congested areas. 
Gay-Lussac (ga-lu-sak') ascended over Paris in a hydrogen-filled 
balloon to a height of 4 miles to get samples of air. Much more re- 
cently, rockets carrying self-sealing bottles brought down air from 
heights as much as 36 miles. 

Although these analyses showed that all the samples of air varied 
only slightly in composition, enough difference was noted to indicate 
that the composition of air is not constant. Hence, air could not be 
a compound. Air, they decided, is a mixture of gases. 

Some other proofs that air is a mixture. The conclusion reached 
by the early investigators that air is a mixture was strengthened by 
further evidence. For example, rain water contains air that has been 
dissolved from the atmosphere. By boiling rain water, this air can be 
driven out and collected. Analysis shows that this air contains almost 
twice as much oxygen as common air. 

If air were a compound, dissolving it in water would not cause any 
change in its composition. But because air is a mixture, each sub- 
stance of which it is composed dissolves in water in proportion to its 
own solubility therein. Since oxygen is more soluble in water than 
nitrogen, the other main component of air, dissolved air contains 
more oxygen than does common air. 





9 lass , -, 
cylinder ',. 

rise of water 

Fig. 18. Finding the per- 
centage of oxygen in the 
air by the use of a eu- 

iron filings 

"','V : '^*?V , 

rise of water 

When scientists succeeded in changing air into a liquid by cooling 
it to a temperature of about 190C., they found that liquid air 
does not have a definite boiling point. Instead of all the liquid air 
boiling at a definite temperature, as does water or any other pure 
compound, they found that nitrogen boiled off first. Other com- 
ponents of the air boiled off at higher temperatures. A pure com- 
pound has a definite boiling point, but liquid air does not. This is 
further proof that air is a mixture. 

Finding the percentage of oxygen in air. The percentage of oxy- 
gen in air is found by using a chemical which will react with oxygen 
and remove it from the air. For example, we might place a small 
piece of white phosphorus on the coiled end of a copper wire and 
insert it in a eudiometer, or measuring tube, which is then inverted 
in a cylinder of water (see Fig. 18) . 

The phosphorus, which is easily oxidized, soon combines with the 
oxygen of the air in the eudiometer according to the equation: 

This reaction continues until the last trace of oxygen has combined, 
forming phosphorus trioxide, white feathery crystals which readily 
dissolve in the water in the cylinder. Phosphorus pentoxide, P 2 O 5 , is 
formed also according to the equation: 

This is a white solid which readily combines with water and, hence, 
is used as a drying agent, especially for gases. 

As the oxygen unites with the phosphorus, a partial vacuum is 
formed in the eudiometer, and the greater air pressure outside the 

* Also written, respectively, P 4 O C and P 4 O 10 . 



tube forces water up the tube. At the end of a few hours the chemical 
action ceases and the water stops rising. The volume of water that has 
risen in the eudiometer is equal to the volume of oxygen that was 
originally present in the air in the measuring tube. Because this rise 
in the height of the water is approximately one-fifth the height of the 
tube, this rise indicates that about 20 percent of air is pure oxygen. 
In performing this experiment, do not touch the phosphorus. 

Instead of phosphorus, a less active element such as iron may be 
used. To perform this experiment, the inside of a eudiometer is 
moistened with water and enough iron filings are added to form 
a thin layer of iron on the inside walls of the tube, as shown in 
Fig. 18. The tube is then inverted in water, and left undisturbed 
overnight. Upon examination, it is found that most of the iron has 
changed to brown rust, an oxide of iron, and the water has risen 
about one-fifth the height of the eudiometer. 

As some molten metals 
harden in the mold, tiny 
holes are formed due to 
the presence of dissolved 
gases. These holes make 
the casting unfit for use. 

To prevent formation of 
holes in the casting, 
nitrogen is bubbled through 
the molten metal in the 
mold. The nitrogen carries 
away the undesirable 
gases. This process is 
known as n/frogen degas- 

Carbide and Carbon Chemicals Company 


Nitrogen is the most abundant element in the atmosphere. Nitro- 
gen is the principal gas that remains after oxygen has been removed 
from air. It was first carefully studied in 1772, by Daniel Rutherford, 
uncle of Sir Walter Scott. Pure nitrogen may be prepared in the 
laboratory by gently heating ammonium nitrite, NH 4 NO 2 . 

NH 4 NO 2 - N 2 T + 2H 2 O 

The physical properties of nitrogen resemble those of its partner, 
oxygen, quite closely. It is colorless, odorless, and scarcely soluble 
in water (about two liters of nitrogen dissolve in 100 liters of water 
under standard conditions) . It is even more difficult to liquefy than 
oxygen, requiring a temperature 13C. lower. It is slightly lighter 
than oxygen. Dry air is about 78 percent nitrogen by volume. 

Chemical properties of nitrogen. Since nitrogen under normal 
conditions is not chemically active, Lavoisier named this element 
azote, meaning devoid of life. However, because it is one of the ele- 
ments found in niter (potassium nitrate) , this name was later 
changed to nitrogen. 

Chemically, nitrogen differs completely from oxygen. It does not 
burn under normal conditions, and does not support respiration. 
It unites with oxygen only at such temperatures as exist in the electric 
arc. However, at higher and carefully controlled temperatures and 
pressures, nitrogen combines with oxygen, hydrogen, and carbon. 
With calcium, magnesium, lithium, and wolfram, it forms a group 
of compounds called nitrides. 

N 2 + O 2 - 2NO (nitric oxide) 

N 2 + 3H 2 2NH 3 (ammonia) 

N 2 + 2C > C 2 N 2 (cyanogen a colorless, poisonous gas) 

N 2 4- 3Mg > Mg 3 N 2 (magnesium nitride) 

Life as we know it consists of plant and animal forms that have 
survived because they were well suited to living in the earth's at- 
mosphere. Life on earth might be radically different if the oxygen of 
the air were not diluted with inactive nitrogen. Nitrogen tones down 
the chemical activity of the oxygen of the air. Hence burning and 
other oxidations are not as rapid as they otherwise would be. 

How nitrogen is used. The vast storehouse of free nitrogen in air 
furnishes an almost limitless supply of this valuable element. At- 
mospheric nitrogen obtained from liquid air is used in making am- 
monia, ammonium compounds, and nitrates. These nitrogen com- 
pounds are essential in the manufacture of explosives and in 
fertilizers (see pages 268, 270, 468, and 470). Nitrogen thus plays 
a dual role, aiding both in supporting life and in destroying it. 


Rayleigh is confronted with another riddle of the atmosphere. In 

1894, Rayleigh (ra'li) , an English scientist, found that "pure" nitro- 
gen obtained from air weighed a little more than an equal volume of 
nitrogen obtained from pure nitrogen compounds. This puzzled him. 
Dalton had declared that the weight of the atom never changed, 
regardless of its source. 

The difference in weight that Rayleigh noticed was very small, and 
might have been ignored as caused by experimental errors. But after 
Rayleigh had spent months investigating this problem, he became 
convinced that the presence of some other element in the air was 
responsible for the difference in the weights of the samples of nitro- 
gen. With the persistence of a true research scientist he finally tracked 
down the cause of this difference. 

The discovery of argon, an inert gas of the atmosphere. Rayleigh, 
together with William Ramsay, studied the experiments of Cav- 
endish and came across the following statement: "I made an experi- 
ment to determine whether the whole of the nitrogen of the at- 
mosphere could be changed to nitric acid. Having condensed as 
much as I could, only a small bubble of air remained. So that if 
there is any part of the nitrogen of our atmosphere which differs 
from the rest, we may safely conclude that it is not more than Y^th 
part of the whole." 

Here was a clue to their problem. Small as this quantity was, 
Cavendish had not treated it as negligible or as an error in his work. 
Rayleigh and Ramsay therefore repeated his experiments and isolated 
a small volume of this gas from the nitrogen of the air. After sub- 
jecting it to every known test, they finally identified a new element, 
heavier than nitrogen, which, because of its chemical inertness, they 
named argon, meaning lazy. 

Besides argon, minute quantities of five other inert gases are 
now known to be present in the atmosphere helium (sun) , neon 
(new) , krypton (hidden) , xenon (strange) , and radon (from ra- 
dium) . The discovery and isolation of these gases from the air are 
other amazing examples of precise and painstaking research. When 
we consider that each of these gases is colorless, odorless, insoluble 
in water, and chemically inert, refusing to unite with even the most 
active elements, we begin to realize why they eluded chemists so long. 

The inert gases go to work. For many years the inert gases re- 
mained chemical curiosities. Helium, first identified in the sun in 
1868 and later found in considerable quantities in natural gas, was 
the first to be put to practical uses. It has taken the place of hydtogen 
in inflating the balloons and blimps of the Army and Navy and some 



of the weather observation balloons of the Weather Bureau. When 
mixed with oxygen, it forms a synthetic air that is used under pres- 
sure in caissons and is supplied to deep-sea divers to prevent the 

The bends are a type of severe cramps caused by the sudden ex- 
pansion and liberation of large quantities of nitrogen gas that have 
entered the blood under the great pressures to which deep-sea divers 
are subjected. When a diver rises to the surface where the air pres- 
sure is considerably lower, less helium is capable of remaining dis- 
solved in his blood, and, therefore, some of this gas is liberated from 
the blood. Since helium is 40 percent less soluble in blood than is 
nitrogen, less helium gas will be forced out of solution by the decrease 
in pressure, and so the substitution of helium for the nitrogen of 
common air helps to prevent the bends. 

Laziness preferred. With the demand for lightweight metal parts 
for airplanes, brought on by the tremendous need for aircraft in 
World War II, research on the problem of welding magnesium was 
stimulated and the helium-atmosphere process was perfected. In an 
atmosphere of helium, oxidation of the magnesium cannot take 
place, and the magnesium or magnesium alloys may be welded with 
ease. The development of this method of welding magnesium is an 

Linde Air Products Company 

This workman is welding structural 
aluminum by the sigma method. The 
term tigma is derived from the first 
letters of the words s/tiefaW inert gas 
metal arc. 


example of the way in which the needs of society and the research 
of scientists are related. 

Argon has replaced helium-arc welding for magnesium, and other 
metals. It is also used to fill electric-light bulbs. When an evacuated 
bulb is in use, metal evaporates from the filament, forming a deposit 
on the inside of the bulb. This deposit blackens the bulb, making 
it very inefficient. In an argon-filled bulb, this evaporation is re- 
tarded and the lamp may be operated at higher temperatures than 
if it had been evacuated only. 

The ruby glow. The inert gases are also widely used in the glowing 
glass tubes so familiar in advertising signs. When an electric current 
is sent through a tube from which air has been removed and a minute 
amount of neon gas introduced, the gas glows with an orange-red 
light. The gas is at low pressure, about 12 mm. of mercury. The 
amount of current required is extremely low, about ^ of an ampere, 
but the voltage varies between 6000 and 12,000 volts. Neon and 
krypton lights are also used to mark airplane routes and to signal 
to airplane pilots. Small neon glow lamps are used in testing high- 
frequency electric circuits such as those in radios. 

When argon is used instead of neon, the light produced is blue. 
However, most of the blue tubes of this type are filled with mercury 
vapor rather than argon. Xenon gives a light blue light, and helium 
a cream-colored or pale orange light. Following page 382 is an il- 
lustration showing the colors produced by the inert gases in lighting. 

The "idle" gases have thus been set to work. A leader in this field 
predicted that, in time, much of our lighting would be done by 
glowing gas in luminous tubes instead of by incandescent filaments. 
Certainly it is true that much outdoor- lighting is now produced in 
this way, but the prediction would not apply to indoor lighting in 
which the field seems definitely being taken over by a newer develop- 
ment, fluorescent tubes (see page 448) . 

Radon gas, enclosed in sealed tubes, is used in the treatment of 
cancer. The exact function of the inert gases in the air is still not 

Our atmosphere also contains water vapor. Water vapor is always 
present in air in varying amounts. The waters of the earth are con- 
stantly evaporating. Plants give off immense quantities of water 
vapor during transpiration, and animals, too, exhale water vapor. 

Rain, dew, snow, fog, and other similar phenomena are caused by 
the condensation of the invisible water vapor of the air. Frequently, 
a pitcher of ice water sweats on the outside. This sweat is the water 
that was formed when the water vapor of the air came in contact 


with the cold outer surface of the pitcher and condensed. To de- 
termine accurately the amount of water vapor in the air, we can 
pass a known volume of air through a drying agent, such as calcium 
chloride, CaCl 2 , or phosphorus pentoxide, P 2 O 5 . The increase in 
weight in the drying agent equals the weight of the water vapor in 
the sample. 

CaCl 2 + 2H 2 O -> CaCl 2 2H 2 O 

Carbon dioxide, too, is present in the air. Millions of tons of 
carbon dioxide are poured into the air daily by the burning of 
organic substances, by the decay of dead plant and animal matter, 
and by the breathing of living things. 

C + O 2 -CO 2 

Carbon dioxide is a colorless, odorless gas fairly soluble in water, 
and 1| times as heavy as air. Because carbon dioxide is already 
completely oxidized, it does not burn. When passed through lime- 
water (a water solution of calcium hydroxide) , it forms a white 
precipitate, calcium carbonate, CaCO 3 . This formation of a white pre- 
cipitate is the common test for carbon dioxide. 

Ca(OH) 2 + CO 2 - CaCO 3 j + H 2 O 

To determine accurately the amount of carbon dioxide present in 
air, a known volume of air is passed through a concentrated solution 
of potassium hydroxide, KOH, and the amount of potassium car- 
bonate, K 2 CO 3 , that forms is determined. 

2KOH + CO 2 - K 2 CO 3 + H 2 O 

Carbon dioxide and its uses are discussed in Chapter 23. 

Cause and effect of atmospheric pollution. An adult inhales about 
37 pounds of air a day, which is five times the weight of the food and 
water that he consumes. We are very careful about getting pure 
water, and have laws to protect us against the sale of impure food. 
But we have not done as much about polluted air. In some cities 
legislation has been passed to cut down pollution of the air caused 
by smoke. Excessive smoke may impair health, damage crops, slowly 
destroy property, and reduce visibility. Smog, a combination of 
smoke and fog, is another serious problem (see pages 329, 616, and 

What is air conditioning? Not so many years ago, it was supposed 
that the air in a crowded room was unhealthful because it contained 


a large percentage of carbon dioxide and a lowered percentage of 
oxygen. It has since been proved that even in a very crowded room 
the percentage of carbon dioxide never reaches a point where it be- 
comes harmful. The amount of oxygen rarely gets below 20 percent, 
and it can be cut down even to 17 percent, at which point a candle 
is extinguished, without being injurious to health. 

Research has shown that "bad air" is really caused by high temper- 
ature, lack of circulation, high percentage of water vapor, and various 
odors that have accumulated. 

Warm air can hold more water vapor than an equal volume of cold 
air. A cubic meter of air at 20C. (68F.) , for example, is capable 
of holding about 17 grams of water vapor, while the same volume of 
air at 11C. (52F.) can hold only about one-half as much. The 
ratio of the weight of water vapor present in air to the weight of 
water vapor it is capable of holding under the existing conditions of 
temperature and pressure is known' as its relative humidity. 

The temperature of the human body is controlled in part by the 
evaporation of perspiration. Evaporation absorbs heat. The amount 
of heat absorbed depends upon the amount of perspiration evap- 
orated. The cooling sensation produced depends upon both the 
amount and the rate of evaporation. 

Air with a high relative humidity, regardless of temperature, will 
evaporate little perspiration. Consequently, on hot, damp days we 
feel hotter than we do on hot days ihat are somewhat drier. 

Tests have shown that in winter the air in many homes has a very 
low relative humidity. Such very dry air has a twofold bad effect. 
It tends to "dry us out." Evaporation of perspiration present proceeds 
rapidly with great cooling effect. Consequently, the temperature of 
dry air must be greater than the temperature of damper air for our 
sensations of warmth to be the same. It t^kes more fuel to maintain 
the higher temperature and, as a result, costs more. In addition, 
temperatures above 21C. (70F.) are likely to produce drowsiness 
and prevent us from doing our best work. Furthermore, such dry air 
tends to dry out the linings of the nose, mouth, and throat, thus 
causing great discomfort and reducing our resistance to common 
colds and other respiratory diseases. 

Today, steady progress is being made in supplying properly con- 
ditioned air, not only to large auditoriums, factories, classrooms, 
banks, office buildings, and railroad trains, but even to subway cars 
and private homes. Air-conditioning equipment filters out dust and 
pollen from the air, exhausts stale air, keeps the relative humidity 
at the right point (about 50 percent) , and maintains a comfortable 







temperature (about 68 F.) . It also destroys the dulling quality of 
"dead" air by keeping the air in motion. In addition to increasing 
mental and physical efficiency, air conditioning is essential in certain 
manufacturing processes, such as printing and the making of pills, 
chocolate, rayon, paper, tobacco products, and steel bearings. 

Summary: The composition of the atmosphere. The chief con- 
stituents of samples of dry air taken near the surface of the earth are 
shown in the following table. 





, Oxygen 
Carbon dioxide 



Air contains variable quantities of water vapor and carbon dioxide; 
small amounts of the inert gases, argon, helium, krypton, neon, 
radon, and xenon; and also minute amounts of other gases, such as 
methane, carbon monoxide, hydrogen, nitrogen dioxide, and ozone, 
as well as very finely divided solids, such as dust, bacteria, spores, and 

Liquid air. In Gullivers Travels, a famous Academy was visited, 
and Jonathan Swift reports how some of its scientists were con- 
densing air and letting the liquid flow like water. Probably Swift 
believed the liquefaction of air a dream never to be realized. Yet 
today liquid air is a common article of commerce and thousands of 
tons of it are used every year. 

The principles underling the manufacture of liquid air are: 
(1) when a liquid evaporates, it absorbs heat from its surroundings 
and thereby lowers their temperature; and (2) the sudden expansion 
of a gas produces this same effect, When air is compressed, cooled, 
and suddenly allowed to expand through a narrow opehing, its 
temperature is lowered. If this process is repeated, a temperature is 
finally reached which is low enough to liquefy the gas. A more de- 
tailed treatment of the liquefaction of gases will be found on 
page 253. 

How the kinetic theory of matter explains the liquefaction of 
air. All matter is thought to be made up of small particles (atoms 
and molecules) that are in constant motion. In the case of gases, this 
motion is extremely rapid (air molecules under normal conditions 



move at about 20 miles per minute) . When any gas is cooled, its 
molecules move more slowly, until finally a temperature is reached 
at which the motion of the molecules of gas is so slow that they come 
close enough together to form larger groups or clusters of molecules, 
thus forming a liquid. When a gas expands, some heat is used in sep- 
arating the molecules. 

Compressing a gas has the effect of bringing the molecules closer 
together. By a simultaneous cooling and compressing, any gas may 
be liquefied. Some gases, such as chlorine, sulfur dioxide, and am- 
monia, are easily liquefied. Other gases, such as oxygen, nitrogen, 
hydrogen, and helium, require extremely low temperatures and 
high pressures to change them to liquids. It has been calculated that 
at 273C. (absolute zero) all motion of the particles of matter 
ceases. The nearest approach to this temperature thus far attained 
was made in 1952 when scientists at our National Bureau of Stand- 
ards reached a temperature within 0.001,5 of absolute zero.* 

* The measurement of absolute temperatures is discussed on p. 644. 

Arthur D. Little, Inc. 

Filling a flask with liquid helium 
in an experimental laboratory. 
Helium is liquefied by subjec- 
ting it to both cooling and pres- 




19. A thermos bottle 
a Dewar flask. Note 
imilarities. What is the 
of the silvered sur- 

Properties and uses of liquid air. Liquid air is a pale blue liquid 
almost as heavy as water. It contains about 21 percent oxygen and 
boils at 190C. When it evaporates, nitrogen boils off first and the 
mixture becomes richer and richer in oxygen, for the same reason 
that the boiling alcohol-water mixture in your car's radiator loses 
alcohol faster than water, because the boiling point of alcohol is 
lower than that of water. 

Liquid air is used chiefly as a source of oxygen, nitrogen, and the 
inert gases, which boil off at different temperatures. The Nazis used 
liquid oxygen and alcohol to fuel the V-2 rocket bombs which they 
hurled against England in 1945. 

Because of the great tendency of liquid air to evaporate, small 
volumes of it are kept in special containers called Dewar flasks, which 
are similar in construction to the familiar thermos bottle. These 
flasks cannot be tightly stoppered, for any attempt to confine liquid 
air too closely results in explosion. Because of the danger of injury 
both from such explosions and from contact with a substance at such 
an extremely low temperature, persons handling liquid air must use 
great care. 

Air Reduction Company, Inc. 

Inferior of an oxygen-nitrogen 
plant. The workman is standing in 
front of a column in which the two 
gases are obtained from liquid air. 


The properties of substances change when immersed in liquid air. 
Liquid mercury becomes solid enough to be used as a hammer head. 
Rubber turns hard and brittle. The resistance of copper to an elec- 
tric current is decreased 50 times. 


Cady, H. P. "Liquid Air." Journal of Chemical Education, 
June, 1931, pp. 1027-1043. A fascinating account of experi- 
ments with liquid air. 

Kaempffert, Waldemar. Explorations in Science, Chapter 7, 
pp. 102-108 is entitled, "This most excellent canopy, the air." 
The Viking Press, New York, 1953. 

Ramsay, William. The Gases of the Atmosphere, pp. 148- 
181, 234-269. The Macmillan Co., London, 1915. Sir William 
tells of the discovery of argon and discusses other inert gases. 


1. Air is a mixture of gases because (1) it has no definite 
composition, (2) when air dissolves in water, the dissolved 
air contains more oxygen than common air, and (3) liquid air 
does not have a definite boiling point. 

2. Air conditioners cleanse the air, keep the relative humid- 
ity where it belongs (about 50 percent) , maintain a comfort- 
able temperature (68 F.) , and keep the air in motion. 

3. The kinetic theory of matter assumes that gases are made 
up of small particles (atoms and molecules) in active motion. 
Cooling the gas slows down this motion and brings the par- 
ticles closer together until a liquid is formed. The molecules 
of liquids and solids are also in motion. 


Group A 

1. Who first definitely proved that air has weight? 

2. (a) What pressure does the atmosphere exert at sea 
level? (b) What causes this pressure? 

3. What are the differences between the troposphere, strat- 
osphere, and Heaviside layer? 

4. Does the composition of the atmosphere prove that air 
is a mixture? Explain. 

5. Give two reasons other than composition for believing 
that air is not a compound. 


6. State the composition of dry air by volume. 

7. What substances are found in air in variable quantities? 

8. Make a diagram illustrating a laboratory method for 
determining the percentage of O 2 in air. 

9. (a) What is the gas that is left in the measuring tube 
used in determining the percentage of O in air? (b) What 
are the impurities in this gas? 

10. State four physical properties of N 2 . 

11. What is the chief chemical property of N 2 ? 

12. Why is it wrong to call N 2 an inert gas? 

13. Write a chemical equation illustrating the action of 
hot Mg on N 2 . 

14. Under what conditions does N., combine chemically 
with O 3 and H 2 ? 

1 I 

15. (a) State two functions of the N 2 of air. (b) State two 
commercial uses of N 2 . 

16. (a) Why is Mg difficult to weld? (b) Under what con- 
dition is it easily welded? (c) Development of this process 
is an example of what? 

17. Name the six inert gases of the atmosphere. 

18. What five properties are common to all six of the inert 
gases of the atmosphere? 

19. Match each element listed in the first column with the 
correct item in the second column. 

1) Ne a) hidden 

2) A b) balloons 

3) Kr c) strange 

4) He d) luminous tubes 

5) Xe e) electric-light bulbs 

f) airplanes 

20. Explain the "sweating" of a pitcher filled with ice water, 

21. How would you determine the amount of water vapor 
present in air? 

22. How would you determine the amount of CO, present 
in air? 

23. State (a) four physical and (b) two chemical properties 
of C0 2 . 

24. What is air conditioning? 

25. Name (a) two causes and (b) three effects of atmos- 
pheric pollution. 


26. What two conditions in the air of a poorly ventilated 
room make it both uncomfortable and unhealthful? 

t . . . 

27. What is relative humidity? 

28. Upon what principle does the liquefaction of a gas 

29. Discuss the meaning of the kinetic theory of matter. 

30. What is the chiet use of liquid air? 

31. State three properties of liquid air. 

32. Compare the effectiveness of determining the percentage 
of O 2 in air by the phosphorus and the iron filings methods. 

Group B 

33. Why does the composition of air by weight differ from 
its composition by volume? 

34. A piece of burning charcoal is plunged into liquid air. 
It keeps on burning with even greater splendor. Why does not 
the extreme cold of liquid air extinguish the burning charcoal? 

35. Explain the presence of vast amounts of free N 2 in air, 
and only relatively small amounts of nitrogen compounds in 
the earth's crust. 

36. In different samples of air the following substances are 
placed: CaCl,, P, hot Mg, and Ca (OH) ,. Explain what hap- 
pens in each case. 

37. Why does a kettle of liquid air boil when placed on ice? 

38. "Coal burned in our furnace returns to us in our 
bread." Explain. 

39. What weight of dry air would be theoretically needed 
to extract ten grams of pure oxygen? 


1. Visit a neon-sign factory or an air-conditioned building 
or factory and make a full report on your visit to the class. 

2. Prepare simple demonstrations to show that (1) air 
contains water vapor, and (2) your breath contains CO 2 . Per- 
form these experiments before your class the following day. 

3. With the aid of a wet and dry bulb thermometer, de- 
termine the relative humidity of your classroom. 

4. Prepare a report on the probable origin of the earth's 
atmosphere. Use a good book on physiography or one of the 
references listed. 





. . . Here evidently we are at the 
birthplace of the chemical equation, 
yet we cannot find in the writings of 
Lavoisier this instrument as we know 
it; for our chemical equation de- 
pends as much on the atomic theory 
as on the doctrine of the conservation 
of mass. Caven and Cranston, 1928 

Chemical equations, the shorthand of chemistry. Chemists use 
symbols and formulas that are understood to represent elements and 
compounds by scientists in all parts of the world. But chemists make 
further use of these symbols and formulas. They use them to tell the 
story of chemical change the reacting substances, the type of chem- 
ical change, the products of the reaction, and various other facts. 
All of these facts are expressed in the form of a chemical equation, 
which, in certain respects, is similar to the equations you have used 
in arithmetic and algebra. 

Considering the chemical reaction that takes place when iron and 
sulfur are heated together, chemists and students of chemistry write 
as follows: 

Fc + S - FcS 

As you see, the two sides of the equation are separated by an arrow, 
-. This arrow means yields, or forms, and should be read using 
these or similar words. On the left side of the arrow, chemists write 
the symbols or formulas of the substances that react. On the right side 
of the arrow, they write the symbols and formulas of the products of 
the reaction. A word-equation that will express the same reaction 
as the chemical equation that we have just been discussing follows: 




Iron reacts with sulfur to yield iron sulfide. 

Equations represent reality. To write a chemical equation, we 
must know the formulas of the substances involved. To write a 
chemical equation correctly, we must know exactly what substances 
are reacting and what substances are the products of the reaction. 
The facts are determined before an equation is written. 

We can write equations for only those chemical reactions known to 
be capable of actually taking place. Chemical equations must repre- 
sent actual conditions. For example, it is incorrect to write He + S 
- HeS, because helium and sulfur do not react together, and 
helium sulfide, HeS, has never been prepared. 

Balancing an equation. Lavoisier once wrote: "One may take it 
-for granted that in every reaction there is an equal quantity of matter 
before and after the operation." Atoms do not disappear in the 
process of chemical reaction. Therefore, the number of atoms in 
the reacting substances must equal the number of atoms in the 
products of the reaction. The number of atoms of each element must 
be the same on each side of the arrow. Every correct equation con- 
forms, in this way, to the law of the conservation of matter. 

Points to remember in balancing equations. 

1) Elements can occur in a free state, not combined with other 
elements. For reasons which you will learn later (pages 279280) , 
a number of these elements, including oxygen, hydrogen, nitrogen, 
chlorine, bromine, and iodine occur as molecules containing 2 atoms. 
They are written O,, H.,. N 2 , C1 2 , Br 2 , and I 2 , thus: 

H 2 + C1 2 2HC1 (hydrogen chloride) 
N 2 4- O 2 > 2NO (nitric oxide) 

. . < Cu 4- Br 2 > CuBr 2 (copper bromide) 

2H 2 O -> 2H 2 + 2 

Standard Oil Company 

Delicate measurements must be 
made to determine the actual 
conditions that are described 
by chemical equations. 


2) In general, a radical remains unaltered during a chemical 
change and its symbols are carried over to the right side of the equa- 
tion unchanged: 

Zn + H 2 SO 4 -> ZnSO 4 + H 2 
2Na + 2HOH - 2NaOH + H 2 
BaO 2 + H 2 SO 4 -> BaSO 4 + H 2 O 2 

Balancing an equation consists of changing coefficients until the same 
total number of atoms of each element is shown on each side of the 

Do not alter the subscript of a radical nor any other subscript in 
order to make an equation balance. Such alteration would mean a 
change in the actual composition of the compound. This in turn 
would mean an entirely different substance which does not actually 
appear in the chemical change. We may, however, alter coefficients 
of any element or compound without changing the meaning of the 
chemical symbols. The coefficient describes the number of molecules 
of a substance. By changing the number of molecules present in a 
chemical change, we do not alter the composition of the substances 
which are involved. 

Your best preparation for writing equations correctly is a thorough 
knowledge of valence and of formulas. In addition, you should 
attack the problem in a systematic and thoughtful way. The follow- 
ing procedure should be helpful. 


1) Write the equation without giving coefficients to any of 
the formulas. 

2) Write any free element occurring in the equation without 
a subscript. Retain all other subscripts. 

3) Select the compound with the greatest number of atoms. 
For one of its elements, compare the number of atoms appear- 
ing on each side of the equation. If the numbers differ, decide 
upon a coefficient or coefficients which will equalize the num- 
ber of atoms on each side. 

4) Repeat this procedure for the other elements until the 
equation is balanced, that is, until the same number of atoms 
of each element appears on each side of the equation. 

5) Add the subscript to any free element in the equation 
which may require it. At this stage, free elements may have 
been given coefficients. If the coefficient of the free element is 


an even number divide the coefficient by two when the sub- 
script is added. The total number of atoms of the free element 
will then remain the same (for example, 6H contains the same 
number of atoms as 3H 2 ) . 

6) If, however, the coefficient of the free element is an odd 
number, it cannot be divided by two. In this case, when the 
subscript is added to the free element, all other formulas on 
both sides of the equation must be multiplied by two. 

EXAMPLE A: Write the equation for the reaction which takes 
place when potassium chlorate is heated to form potassium 
chloride and oxygen. 

1) Write the equation without coefficients. Note that the 
free oxygen which is usually written O 2 is written at this stage 
without a subscript. 

KClOa -^ KC1 + O 

2) KC1O 3 is the compound with the largest number of 
atoms. By inspection, we see that oxygen is the only element 
in the compound which does not appear with an equal number 
of atoms on the other side of the arrow. To make the O on the 
right balance the O 3 on the left, we multiply O by the coeffi- 
cient 3: 

KC1O 3 -> KC1 + 3O 

3) As we have just learned, free oxygen must be written as 
O.,. Add the subscript to the free oxygen on the right of the 
arrow. The coefficient 3 is not divisible by two, so when we 
add the subscript, we find there are six atoms of free oxygen 
on the right side of the arrow, but only three atoms of oxygen 
on the left side: 

KC1O 3 -> KC1 + 3O 2 

4) We may make the equation balance again by multiplying 
.all other formulas by two: 

2KC1O 3 -> 2KC1 + 3O 2 

EXAMPLE B: Write the equation for the burning of benzene 
(C 6 H ) to form CO 2 and H 2 O. 

1) Write the unbalanced equation. Do not write the sub- 
.script for the free oxygen. 

O -> CO 2 + H 2 O 


2) C 6 H 6 is the compound with the largest number of atoms. 
It contains 6 atoms of carbon, but there is only one atom of 
carbon to the right of the arrow. Therefore, we place the 
coefficient 6 in front of CO. 

3) C 6 H 6 also contains 6 atoms of hydrogen, but there are 
only 2 atoms of hydrogen on the right side of the arrow. There- 
fore, we place the coefficient 3 in front of H 2 O. 

CeH 6 + O -> 6CO 2 + 3H 2 O 

4) We now see that there are 15 atoms (6 X 2 + 3 X *) of 
oxygen on the right, but only one atom of oxygen on the left. 
Therefore, we place the coefficient 15 in front of O. 

C 6 H fl + 150 -> 6C0 2 + 3H 2 O 

5) Since free oxygen is written O 2 , we add the subscript. 

6) Since 15 is not divisible by two, we may bring the equa- 
tion into balance once again by multiplying all other formulas 
by the coefficient 2. 

SQHe + 15O 2 -> 12CO 2 + 6H 2 O 

If the coefficient of the oxygen had been an even number, we 
could have divided it by two when adding the subscript. It 
would then be unnecessary to multiply the other formulas by 
the coefficient 2. 

HOW TO CHECK EQUATIONS. To check an equation, proceed as follows: 

1) Examine the first symbol on the left. Compare the num- 
ber of atoms against the number of atoms of the same element 
on the other side of the equation. If the numbers are equal, 
put a small check over the symbols for the element. 

2) Proceed to each symbol in turn until you have put a check 
over every symbol in the equation. 

Thus, in checking the final equation in Example A above, 
2KC1O 8 -* 2KC1 + 3O 2 , examine the first symbol, K. There are two 
atoms of potassium on each side. Put a check over the K on each side. 
Examine the next symbol, Cl. Again we have two atoms on each 
side. Check them. Finally, we have six atoms of oxygen on each side, 
which we check. All the symbols are then checked and accounted for: 

v/ v/ v/ V \/ v/ 

2KC1O 3 -> 2KC1 + 3O 2 



The equation is now balanced accurately. It is correct both mathe- 
matically and chemically. Following the same procedure, check the 
final equation in Example B. 

Fig. 20. The of th tuft contains o 

of oncf an test holding a 

of AgN0 3 . When the | t inverted 

(right), ACI is but the cca | e t ho 

no change. No matter hot 





To be sure that you understand the suggestions just given for bal- 
ancing equations, balance the following: 

a) NaClOs -> NaCl + O g) 

b) Hg + 0-HgO h) 

c) P 2 5 + H 2 -> HPO, i) 

d) Mg + 0->MgO j) 

e) Na + Cl->NaCl k) 

f) C + O->CO 

Gu + O -> CuO 
Fe -f O > Fe 2 O 3 
O + P - P 2 
P 4- O -> P 2 O 5 
SiO 2 -f G -* SiC 


What are the four types of chemical reactions? There are four 
general types of chemical reactions. 

1) Direct combination, or synthesis. When two or more elements 
or compounds combine directly, forming a chemically more complex 


t . . . 

6. Why must subscripts never be changed in balancing 

7. What is wrong with the following: 

2Ne (neon) + O 2 - 2NeO? 

8. Why may coefficients be changed in balancing an equa- 

9. Why is a knowledge of valence so essential in writing 

10. Name five elements which, when free, must be written 
with the subscript 2. ft 

a 1 

11. Correct this: Mg -(- Br - MgBr 2 . 

12. Generally, what happens to a radical during a chemical 

13. Balance the following equation, giving your reason for 
each step: 

Acetylene (C 2 H 2 ) -f oxygen (O 2 ) 

-> carbon dioxide (CO 2 ) -\- water (H 2 O) 

14. What are the four general types of reactions? 

15. Give an example of an equation illustrating direct 

16. Give an example of an equation representing simple 

17. How do you recognize an equation representing simple 

18. Give an example of an equation representing double 

19. What is the general rule followed in balancing a double- 
replacement equation? 

20. Complete and balance the following equation. Show 
how to check it to see if it is correctly balanced. 

Na -f HOH - NaOH + H 

21. Show in detail how to balance the equation represent- 
ing the decomposition of sodium chlorate, NaClO 3 . 

22. When CaCl 2 -2H 2 O is heated, water is liberated. Write 
the equation. 


23. Balance the following equations. Check each one. Do 
not write in this book. 

a) Cu + S->Cu 2 S 

b) P + C1 2 ->PC1 3 

c) C + C0 2 ->CO 

d) H 2 S + Pb(NO 3 ) 2 -> HN0 3 + PbS 

e) CaO + G -> GaC 2 + CO 

f) SiO 2 + G - SiC + CO 

24. Complete and balance the following equations. Check 
each one. Do not write in this book. 

a) H 2 O 2 > e) BaO 2 + H 2 SO 4 > 

b) HgO-> f) AgNO 3 + KCl- 

c) Fe + HCl-> g) Mg + H 2 S0 4 - 

d) CuO + H 2 - h) Ca(OH) 2 + CO 2 -> 

25. Write balanced equations for the following. Name each 
element and compound that appears in the equation. Check 
each equation. 

a) The electrolysis of H 2 O. 

b) The laboratory preparation of O 2 . 

c) The equation representing Priestley's discovery of O L> . 

d) The laboratory preparation of H 2 . 

e) The heating of Cu in air. 

f) The preparation of H 2 O 2 . 

g) The reduction of CuO^ by H 2 . 

h) The heating of crystallized washing soda, 
i) The chemical reaction of the oxy hydrogen torch, 
j) The action of Na on H 2 O. 
k) The passing of N 2 over hot Mg. 
1) The reaction between H 2 O and CO 2 . 

26. Is there anything in a balanced equation which did not 
have to be determined by experimentation? 



. . . As the usefulness and accuracy 
of chemistry depend entirely upon 
the determination of the weights of 
the ingredients and products, too 
much precision cannot be employed 
in this part of the subject, and for 
this purpose we must be provided 
with good instruments. 

Lavoisier, 1743-1794 

Measurement in chemistry. Chemists are chiefly concerned in 
analyzing substances or in making new substances. One question 
they always ask is how much of each element is present in a sub- 
stance. Their measurements of quantity must be very exact. Some- 
times their answer is in terms of the volume occupied by an element. 
More often the answer they want is in terms of weight. 

Here in America, in everyday affairs, we measure the weight of 
substances in ounces and pounds. Chemists use a different kind of 
measurement. They weigh substances in grams, kilograms, and micro- 
grams. They also measure the weight of atoms in terms of atomic 

All measures of weight are comparisons. When you say, "My 
friend, Charlie Ross, weighs 132 pounds," you are making a definite 
statement that anyone can understand. But what is a pound? You 
know what a pound of butter looks like, but a pound by itself 
doesn't look like anything. 

Down in the Bureau of Standards in Washington is a carefully 
protected metal cylinder called the prototype, or standard, kilogram. 
Every measurement of weight made in our nation is a comparison 
with that particular kilogram. For example, the pound which we 



use in everyday measurements of weight may be defined as 0.4536 
of the standard kilogram. 

Just as the pound is divided into ounces, the kilogram may be 
divided into smaller units known as grams. One thousand grams 
equal one kilogram. The gram is a customary unit of weight in chem- 
istry. When chemists speak of ten grams of iron, they are actually 
referring to an amount weighing T ^ of the standard kilogram in 

Chemists also use a Table of Atomic Weights in which one atom of 
each element has a particular weight. The weight assigned to each 
element, like all measures of weight, is a comparison. How was this 
Table of Atomic Weights made? 

Dalton determines atomic weights. As you know, Dal ton believed 
that the atoms of different elements have different weights. Dalton 
knew that if he could find the weights of the atoms, the progress of 
chemistry would be speeded. He rcali/.ed that he could not actually 
weigh an atom of an element. In fact, it took more than 100 years 
from the time of Dalton's experiments until accurate methods and 
precise instruments made it possible to determine the actual weights 
of atoms. 

However, Dalton knew that elements combine according to fixed 
ratios by weight. For example, 22.997 grams of sodium combine with 
79.916 grams of bromine to form sodium bromide a ratio of ap- 
proximately 1 : 3.5. The two elements never vary in the ratio of their 
weights in forming this compound. In a similar manner, all elements 
combine according to certain specific ratios by weight. Dalton be- 
lieved that the ratios of these combining weights depended upon the 
weights of the atoms of each element. He believed further that, by 
studying the ratios of the weights in which elements combine, he 
would be able to determine the relative weights of the atoms of the 

The prototype kilogram No. 20 
is a platinum-iridium cylinder 
39 mm. in diameter and 39 mm. 
high. It is kept at Washington, 
D.C., in the laboratories of the 
Bureau of Standards. 


He selected the atom of hydrogen for his standard and assigned to 
it the atomic weight 1. By choosing the lightest element for a stand- 
ard, he made sure that the atomic weight of each of the other ele- 
ments would be greater than one. Then in a series of experiments, 
he found how the weight of the atoms of 13 other elements com- 
pared with his standard. He found that an atom of oxygen weighed 
seven times as much as an atom of hydrogen and assigned to oxygen 
the atomic weight of 7; he found that an atom of phosphorus 
weighed nine times as much as an atom of hydrogen and assigned to 
phosphorus the atomic weight of 9. These weights have since been 
found in error, but that in no way detracts from the value of Dalton's 

Dalton prepares an historic table. From his experimental data, 
Dal ton prepared a list of the 14 elements arranged according to the 
increasing relative weight of their atoms. This was the first table of 
atomic weights. It was first made public on October 21, 1803 "before 
a select group of nine members and friends in the rooms of the 
Literary and Philosophical Society of Manchester." Although inac- 
curate, the table compiled by this Quaker schoolteacher remains a 
monument to his foresight and intellectual accomplishment. His 
achievement was a crucial advance in chemistry and formed the 
keystone of his theory. 

What is the atomic weight of an element? Later investigators fol- 
lowed Dalton's method, but used oxygen as a standard rather than 
hydrogen. They found oxygen a better choice because it combines 
with far more elements than hydrogen. The atomic weight of oxygen 
was given the whole number 16. Many of the other elements have 
atomic weights that are whole numbers. The weight of hydrogen 
remains approximately one actually 1.0078 in comparison to 

The modern Table of Atomic Weights appears on the opposite 
page. For practical purposes, since the atomic weight of hydrogen is 
approximately one, hydrogen can still be used as the basis of com- 
parison. The number 200.6 after mercury means that one atom of 
mercury weighs a little less than 200 times as much as an atom of 
hydrogen. It weighs a little more than 12| times as much as an atom 
of oxygen. 

The atomic weight of an element is a number that shows the 
comparison of the weight of one of its atoms to the weight of one 
atom of oxygen, which is considered to be 16. 

Always remember that atomic weights are not measured weights 
like an ounce or a gram. They are merely relative weights. The 




AT. WT. 



AT. WT. 



















































































































actual measured weight of each of the atoms is an extremely small 
quantity, inconvenient to use in most calculations. For example, 
the actual weight of an atom of oxygen is about 0.000,000,000,000,- 
000,000,000,026 gram. 

The Table of Atomic Weights is the foundation of chemical 
mathematics. Realizing the importance of accurate atomic weights, 
many chemists performed a tremendous number of experiments to 
make the Table of Atomic Weights as free from error as possible. 
Chemists of all countries cooperated in this huge undertaking. In 
our own country, Theodore W. Richards and his students at Harvard 
spent almost half a century in this epoch-making work. National 
boundaries were forgotten and men from all over the world pooled 
the results of their experiments to give us our present Table of 
Atomic Weights. Research chemists, industrial chemists, and student 
chemists all depend upon the Table of Atomic Weights in making 
their chemical calculations. 

Solving type problems in chemistry. Most of the problems met in 
elementary chemistry can be grouped conveniently into five types. 
The first two, and one variety of the third are described below. Two 
additipnal varieties of the third type are discussed in Chapter 19 
and two more types, more complex and less frequently met, are 
discussed on pages 636641. 



With a thorough understanding of the type problems discussed 
in this book, you should have no trouble in solving practically all 
the common chemical problems. Frequent reflective practice is, of 
course, necessary for mastery. Hence problems of various types are 
included in the questions at the end of each of the remaining 

The meaning of symbols. A knowledge of the meaning of chemical 
symbols is essential in solving the various types of chemical problems. 
The symbol for an element, like K, does three different jobs. First, 
it may be used to name the element. Second, it may be used to mean 
one atom of that element. Third, it may stand for one atomic weight 
of the element. Thus K is the symbol for potassium, for one atom of 
potassium, and for 39 units of potassium in any system of measuring 
weight (as 39 grams, 39 ounces, 39 pounds) . 

The chemical symbols for a formula like KCl also do three jobs. 
First, they represent the name of the compound, potassium chloride. 
Second, they represent one molecule of potassium chloride. Third, 
they represent one molecular weight of the compound. All of the 
type problems involve the use of both atomic iveight and molecular 



Molecular weight is the ratio of the weight of one molecule 
of a compound to the atomic weight of oxygen (16) . Like 
atomic weight, it is only a relative weight. The molecular 
weight of a compound is obtained by adding together the 
atomic weights of each of the atoms in one molecule of the 


1. Find the atomic weights of the elements in the chart on 
page 127. Place these numbers under the symbols and add 

* Some compounds do not exist as molecules and therefore cannot have a 
molecular weight. However, the formulas of these compounds are written in the 
conventional manner. The steps given in this discussion for finding molecular 
weight may also be used for finding the formula weight of the non-molecular 
compounds. For the purposes of this book, formula weight may be considered the 
same as molecular weight. 


EXAMPLE: Find the molecular weight of potassium chloride,. 

K a 

39 + 35.5 = 74.5, the mol. wt. of KC1 

2. When an element is followed by a subscript, be sure to 
multiply the atomic weight of the element by the subscript. 

EXAMPLE: Find the molecular weight of magnesium sulfate, 
MgS0 4 . 

Mg S 4 

24 + 32 + (16 X 4) = 24 + 32 + 64 

= 120, the mol. wt. of MgSO 4 

3. When a formula is preceded by a numerical coefficient, 
multiply the total molecular weight by the coefficient to find 
the relative weight (rel. wt.) . 

EXAMPLE: Find the relative weight of two molecules of mag- 
nesium carbonate, 2MgCO 3 . 

2(Mg C 8 ) 

2 [24 + 12 + (16 X 3)] 

2 (24 + 12 + 48) 

2 (84) = 168, the rel. wt. of 2 molecules of MgCO 3 

4. When a radical is enclosed in parentheses followed by a 
subscript, multiply the sum of the atomic weights of all the 
elements of the radical by the subscript. 

EXAMPLE: Find the molecular weight of calcium bicarbonate, 
Ca (HC0 3 ) 2 . 

Ca (H C O 8 ) 2 

40 + [1 + 12 + (16 X 3)]2 

40 + (1 + 12 + 48)2 

40 +(6 1)2 

40 + 122 = 162, the mol. wt. of Ca (HCO 8 ) 2 

5. Water of crystallization (see page 68) is chemically part 
of certain compounds and is separated from the rest of the 
formula by a dot. The dot stands for plus and is not to be 
considered a multiplication sign. 


EXAMPLE: Find the molecular weight of crystallized copper 
sulfate, CuSO 4 5H 2 O. 

Cu S 4 5H 2 O 

64 + 32 + (16 X 4) + 5(1 X 2 + 16) 
64 + 32 + 64 +5(18) 

64 + 32 + 64 + 90 = 250, the mol. wt. of crystallized copper 


The gram-molecular weight or mole. A very convenient unit in 
many calculations is the gram-molecular weight or mole, which is 
used in chemical equations (see pages 282-283) and in preparing 
standard solutions (see pages 207-208) . A mole is the molecular 
weight of a substance expressed in grams. For example, a mole of 
potassium chloride (see Procedure 1) is 74.5 grams. 


1. Calculate the molecular weight of (a) KBr and (b) Nal. 

2. Find the molecular weights of (a) LiCl and (b) ZnO. 

3. What is the weight of a mole of K 3 PO 4 ? 

4. Find the molecular weights of (a) H 2 SO 4 , (b) CaCO 3> 
and (c) BaSO 3 . 

5. Calculate the molecular weights of (a) Cu (HCO 3 ) , 
(b)Ba(N0 3 ) 2 , and (c) A1 2 (SO 4 ) ,. 

6. Calculate the molecular weight of Na 2 S 2 O 3 5H 2 O (com- 
monly called hypo) . 

7. What is the weight of a mole of gypsum, CaSO 4 2H 2 O? 

8. Find the molecular weight of plaster of Paris, 

(CaS0 4 ) 2 .H 2 0. 

9. Thfe formula of glauber salt is Na 2 SO 4 10H.,O. Find 
the relative weight of three molecules of this substance. 

10. Calculate the relative weight of 4Na 2 B 4 O 7 10H 2 O. 


The percentage composition of a compound is found by 
computing the percentage by weight of each different element 
in the compound. This is a simple percentage problem. 

Procedure. Divide the atomic weight of each element by the molecu- 
lar weight of the compound and multiply the fraction thus 
obtained by 100. 


EXAMPLE A: Find the percentage composition of nitric acid, 
HNO 8 . 

H N O 3 

1-1- 14 + (16X3) 

1 + 14 + 48 = 63, the mol. wt. of HNO 3 
% of hydrogen 

at.wt.ofH 1 X 100 

mol.wtofHNO, ~ 63 

% of nitrogen 

x 100 - 22 . 2% 

mol. wt. of HNO 3 --- 63 

% of oxygen 

_ rel.wt.of30 48X100 


mol. wt. of HN0 3 63 

Total = 100.0% 

EXAMPLE B: Find the percentage of water of crystallization in 
BaCl 2 - 2H 2 O. 

Ba C1 2 2H 2 O 

137 + (35.5 X 2) + 2(1 X 2 + 16) 

137 + 71 +2(18) 

137 + 71 + 36 = 244, the mol. wt. of BaCl 2 2H 2 O. 

Percentage = rcl. wt. of 2H 2 O x 1Q() 
of water mol. wt. of crystal 

EXAMPLE C: Find the percentage composition of 
Na 2 SO 4 10H 2 O. 

Na 2 S O 4 10H 2 O 

(23 X 2) + 32 + (16 X 4) + 10(1 X 2 + 16) 
46 + 32 + 64 + 180 = 322, the mol. wt. of the compound. 

Now find the total atomic weight of each element in the 
compound, thus: 

Sodium - 2 atoms - 23 X 2 = 46, the total at. wt. of Na 
Sulfur = 1 atom = 32, the total at. wt. of S 

Oxygen = 14 atoms = 16 X 14 = 224, the total at. wt. of O 
Hydrogen 20 atoms = 20 X 1 = 20, the total at. wt. of H 
Sum of the at. wt. = 322, the mol. wt. of the 



% of sodium = -3 X 100 = 14.3% 
% of sulfur = ^ X 100 = 10.0% 
% of oxygen - ||| X 100 - 69.5% 

% of hydrogen = ~ X 100 = 6.2% 

Total = 100.0% 

EXAMPLE D: Find the weight of iron in 80 Ib. of an ore con- 
taining 90 percent ferric oxide, Fe 2 O 3 . 

The weight of ferric oxide in the ore is 90 percent of 80 Ib., 
or 72 Ib. 

Fe 2 O 3 

(56 X 2) + (16 X 3) = 160 = mol. wt. of Fe 2 O 3 
Percentage of iron in Fe 2 O 3 = H$ X 100 = 70% 
Therefore 70% of 72 Ib., or 50.4 Ib., is the weight of the iron 
in the ore. 


1. Calculate the percentage composition of (a) water, 
(b) hydrogen peroxide, and (c) mercuric oxide. 

2. Calculate the percentage composition of (a) H 2 CO 3 , 
(b) N 2 O 4 , and (c) CaSO 4 - 2H 2 O. 

3. Find the percentage composition of chrome alum, * 

KCr(S0 4 ) 2 - 12H 2 0. 

4. Calculate the percentage composition of crystallized potas- 
sium ferrocyanide, whose formula is K 4 Fe (CN) 6 3H 2 O. 

5. Find the percentage of oxygen in a compound whose 
formula is NiSO 4 - (NH 4 ) 2 SO 4 - 6H 2 O. 

6. How much aluminum can be obtained from 100 Ib. of its 
cryolite ore which, upon analysis, shows the presence of 80 
percent Na 3 AlF ? 


Because the symbol of an element and the formula of a com- 
pound may represent definite weights, an equation also may 
be considered to represent definite weights of the substances 
taking part in the reaction. Thus, 2Ag -f- S Ag 2 S may be 
read, 216 grams of silver, plus 32 grams of sulfur yield 248 
grams of silver sulfide. Note that the actual weights are based 
on the atomic weights. 

Problems based on chemical equations may be broadly di- 
vided into three groups: A. Straight- weight problems; 
B. Weight-volume problems; and C. Straight- volume problems. 


Group A is discussed below. Groups B and C are discussed 
later in the book. 


Straight-weight problems involve finding one weight in an 
equation when another is given. 

EXAMPLE: How many grams of calcium carbonate will be 
formed by the complete reaction between 222 g. of calcium 
hydroxide and carbon dioxide? 


1. Write the balanced equation. 

Ca(OH) 2 + C0 2 - H 2 + CaCO 3 

2. Write the given weight over its formula. Write x over the 
formula of the substance whose weight is to be found. Cross 
out all other formulas in the equation. 

222 g. x g. 

Ca(OH) 2 +^Q*,^H2QL+ CaCO 8 

3. Since the same relationship exists between the actual 
weights as exists between the molecular weights represented 
in the equation, write the molecular weights of the substances 
involved under their respective formulas. Do not ignore any 

222 g. x g. 

Ca (O H) 2 - Ca C O 3 

40 + (16 + 1)2 40 + 12 + (16 X 3) 

40 + (17X2) 40+ 12 + 48 

74 100 

4. Write the mathematical equation represented by the 
known and unknown weights. Solve for x. 

222 = __ 
74 100 
74* = 22,200 

x = 300, the number of grams of CaCOs produced 

Alternate method. We can avoid the use of an equation in- 
volving x by solving the problem as follows: 


wt. of substance used 


mol. wt. of sub- 
X P - = Answer 

mol. wt. of substance used stance formed 
X 100 = 300, the number of grams of CaCO 3 produced. 


1. How much magnesium is required to react with suffi- 
cient hydrochloric acid to produce 10 g. of hydrogen? 

Mg + 2HC1 -* MgCl 2 + H 2 

2. 434 g. of mercuric oxide, HgO, were decomposed by 
heat. How much mercury was formed? 

3. How much potassium chlorate would be needed to 
prepare 384 g. of oxygen? 

4. How much hydrogen would be needed to reduce com- 
pletely 100 g. of cupric oxide, CuO? 

5. 11.5 g. of sodium react completely with water. What 
weight of sodium hydroxide is formed? 

6. By the electrolysis of water 12 Ib. of hydrogen were 
liberated. What weight of oxygen was formed at the same 

7. (a) What weight of magnesium will be needed to react 
with sulfuric acid to produce 30 g. of MgSO 4 ? (b) What 
weight of hydrogen will be evolved? 

8. After heating, 10 g. of crystalli/ed copper sulfate gave 
6.4 g. of the anhydrous salt, CuSO 4 . Calculate the number of 
molecules of water of crystallization in the original compound. 

Let x represent the number of molecules of water of crystal- 

10 g. 
CuSO 4 -*H 2 O 

6.4 g. 
GuSO 4 

3.6 g. (that is, 10 - 6.4) 

Now complete the problem. 

Standard Oil Company (A'./.) 

Delicate instruments such 
as these are used in micro- 
chemistry, the branch of 
chemistry which involves 
handling extremely small 
quantities of matter. 


9. After being heated, 10 g. of crystallized washing soda 
gave 3.71 g. of Na 2 CO 3 . Calculate the number of molecules 
of water of crystallization. 

10. If 4 g. of crystallized barium chloride lost 0.59 g. upon 
being heated to constant weight, find the formula of the 
crystalline salt. 


Jaffe, Bernard. Chemical Calculations. World Book Co., 
Yonkers, 1947. A systematic presentation of the solution of 
type problems, with 1000 problems arranged progressively 
according to lesson assignments. 

Kendall, James. At Home among the Atoms. D. Appleton- 
Century Co., New York, 1932. "A Few Figures" explains atomic 
weights in a novel way. Tells how atomic weights are found. 


1. The atomic weight of an element is a number that rep- 
resents the ratio of the weight of 1 of its atoms to the weight 
of 1 atom of oxygen. Atomic weights are all relative weights. 

2. The molecular weight of a compound is the ratio ol the 
weight of 1 molecule of the compound to the atomic weight of 
oxygen (16) . 

3. A chemical symbol, in addition to representing an ele- 
ment, represents one atomic weight of that element. 

4. A mole of a substance is its molecular weight expressed 
in grams. 


Group A 

1. (a) Who issued the first table of atomic weights? (b) 
Why was it later decided to use the weight of the oxygen atom 
as a standard instead of the weight of the hydrogen atom? 

2. Exactly what is meant by saying that oxygen has an at. 
wt. of 16? 

3. Find the mol. wt. of the compounds that have the fol- 
lowing formulas: (a) cupric acetate, Cu (C 2 H 3 O 2 ) 2 H 2 O; 
(b) chloroplatinic acid, H 2 PtCl 2 ; (c) microcosmic salt, 
HNaNH 4 PO 4 -4H 2 O. Refer to Table 4 on page 127. 


I ... 

4. Find the percentage composition of each of the follow- 
ing compounds: (a) BaCO 3 , (b) KMnO 4 , (c) K 4 Fe (CN) 6 . 
Check each result. 

5. Determine the percentage composition of each of the 
following compounds: (a) BaSO 4 ; (b) KCN; (c) (NH 4 ) 2 CO,. 

6. Calculate the percentage of H in alum, 

KA1(S0 4 ) 2 -12H 2 0. 

7. Find the percentage of water of crystallization in 

Sr (N0 3 ) 2 5H 2 0. 

8. A ton of limestone, CaCO a , was heated in a lime kiln 
until all of it was changed to quicklime, CaO. The equation 
for this reaction is: CaCO 3 - CaO -+- CO 2 j. How much 
quicklime was formed? To answer this question, first decide 
what type ot problem this is: What four steps have you learned 
to take in solving such a problem? Solve the problem. 

9. Find the weights of (a) H 2 and (b) ZnSO 4 formed by 
the complete reaction of 130 g. of Zn and sufficient H.,SO t . 

Group B 

10. 320 g. of Fe^Oy, on being reduced, form 224 g. of Fe. 
What is the at. wt. of oxygen? 

11. 11.95 g. of lead sulfide, PbS, will produce 10.35 g. of 
lead. From this fact, and knowing that the at. wt. of S is 32, 
calculate the at. wt. of Pb. 

12. Of what use to the manufacturing chemist is knowledge 
of the percentage composition of a compound? Select from 
the chapter on Cu an ore of that metal and illustrate. 

13. Suppose a chemist were going to manufacture HC1 from 
NaCl and H 2 SO 4 . What helpful information could he gain 
from the following equation representing the reaction that 
would occur? 2NaCl + H 2 SO 4 >Na 2 SO 4 + 2HC1 


1. Consult a manufacturing chemist or an analytical chemist 
and discuss with him how he uses chemical mathematics in 
his business or profession. Report your findings to your class. 

2. Construct a large graph to represent the percentage of 
each element in crystallized washing soda, Na 2 CO 3 10H 2 O. 
Use different colors for each element. Show also the per- 
centage of HQ*8JP;this compound. 



. . . Search for the truth is in one 
way hard and another easy, for it is 
evident tliat no one can master it 
; fully nor miss it wholly. But each 

adds a little to our knowledge of Na- 
ture and from all the facts assembled 
there arises a certain grandeur. 


A Swedish druggist discovers chlorine. One of the scientists who 
lived and worked at about the same time as Priestley, Cavendish, and 
Lavoisier was a Swedish pharmacist who is well known to the world 
of chemistry. He not only prepared oxygen earlier than Priestley, 
but also, in 1774, discovered chlorine. Scheele is the only great chem- 
ist whose whole lifework was accomplished behind the counter or 
in the prescription laboratory of one drugstore or another! 

When he had discovered the greenish-yellow chlorine gas that 
made his nose and throat sting and almost blinded him, he wrote to 
a friend: "Oh how happy I am; 1 seldom think of eating, or drinking, 
or where I live; I scarcely pay attention to my pharmaceutical busi- 
ness. But to watch new phenomena, this is my consuming interest." 
At 43, as a result of prolonged exposure to unhealthful conditions in 
his crude laboratory, Scheele died a martyr to the rapidly develop- 
ing and expanding science of experimental chemistry. 

How chlorine is prepared in the laboratory. Hie method used by 
Scheele to prepare chlorine is still the common laboratory method 
used today. Two compounds are used manganese dioxide and 
hydrochloric acid. When this mixure is heated, chlorine is liberated 
from the hydrochloric acid. 



MnO 2 +HCl 

Fig. 21. Laboratory preparation of chlorine. What 
property of chlorine makes it wise to pass the excess 
gas into water? 

Although chlorine gas is fairly soluble in water, it may be col- 
lected by water displacement. Since it is heavier than air, it is us- 
ually collected by displacing air (see Fig. 21) . 

The concentrated hydrochloric acid used in this reaction supplies 
the chlorine. In this reaction the manganese dioxide acts as an oxi- 
dizing agent, combining with the hydrogen of the acid and liberating 
free chlorine. The equation for this reaction is: 

[4HjCl + MnJO^i - 2H 2 O + MnCl 2 + C1 2 

This equation may be considered to represent two reactions. The 
first is of the double-replacement type: 

MnO 2 + 4HC1 - MnCl 4 + 2H 2 O 

and the second, a simple decomposition: *. 

MnCl 4 - MnCl 2 + C1 2 

Other oxidizing agents such as potassium chlorate or lead diox- 
ide, PbO 2 , may be used instead of manganese dioxide. 

The physical properties of chlorine. Chlorine is a greenish-yellow 
gas, two and one-half times as heavy as air. It is fairly soluble in 
water, forming yellowish chlorine water (one volume of water dis- 
solves about two volumes of chlorine gas under normal laboratory 
conditions) . It has a penetrating odor and attacks the membranes of 
the nose, throat, and lungs. Inhaling ammonia or alcohol vapor 
counteracts this irritating action to some extent. Chlorine was the 
first gas to be liquefied. It liquefies at about 34C. at normal 
pressure. Faraday, who was the first to liquefy it, wrote to a friend 
in 1823, "I hope to be able to reduce many other gases to the liquid 
form." He did. 



Chemical properties of chlorine. Chlorine is a typical nonmetal. 
It has a valence of one, that is, it combines with monovalent hydro- 
gen, atom for atom. Chlorine is very active chemically. It unites 
with nearly all other elements, forming compounds called chlorides, 
just as oxygen forms oxides. Thus, a bit of sodium reacts brilliantly 
when heated with moist chlorine. 

Here is a thrilling example of the marvels of chemistry. Sodium, a 
silvery, poisonous solid, unites with chlorine, a greenish-yellow, 
poisonous gas, producing a white solid, common salt, which is essen- 
tial in the diet of both man and all animals. 

2Na + C1 2 -> 2NaCl 

Chlorine has a strong attraction for hydrogen. When hydrogen 
and chlorine are mixed and exposed to strong light or ignited by a 
spark, they combine with explosive violence, forming hydrogen 
chloride gas. The equation for this reaction is: 

H 2 +C1 2 ->2HC1 

If a jet of hydrogen burning in air is thrust into a jar of chlorine, 
it will continue to burn, giving off hydrogen chloride as the product 
of combustion. 

So powerful is the attraction of chlorine for hydrogen that it will 
tear hydrogen away from some of its compounds. Thus, when tur- 
pentine, a compound of carbon and hydrogen, is poured over a piece 
of filter paper which is then thrust into a bottle of chlorine, a flash 
of light occurs and a black powder is formed. The black powder is 
the free carbon which remains after the hydrogen from the turpen- 
tine has combined with the chlorine, forming hydrogen chloride. 
Chlorine will combine with water, liberating oxygen. 

In this industrial pla 
chlorine is being man 
factored in mercury cells. 


How chlorine is prepared for industrial use. The electrolysis of 
brine, a water solution of sodium chloride, is the source of most of 
the chlorine used today. The electric current liberates free chlorine 
at the carbon anode. At the cathode the sodium liberated reacts 
immediately with the water, forming free hydrogen and sodium 

2NaCl -> C1 2 + 2Na 
2Na + 2H 2 O -* 2NaOH + H 2 

The combined equation, then, is: 

2NaCl + 2H 2 O - C1 2 1 + H 2 \ + 2NaOH 

The chlorine gas is dried by passing it through concentrated sul- 
luric acid, and then liquefied. The yellow, liquid chlorine is stored 
in steel cylinders, each holding from one to 300 pounds of chlorine 
free from water vapor. This process, by which three valuable prod- 
ucts chlorine, hydrogen, and sodium hydroxide are formed 
from two low-cost, plentiful compounds by means of an economical 
outlay of electric energy, is described in greater detail on page 212 
under the industrial preparation of sodium hydroxide. Some chlo- 
rine is also obtained by the electrolysis of molten NaCl (see page 

The scientist serves humanity. Soon after the discovery of chlorine, 
Berthollet hit upon the idea of using the bleaching action of chlorine 
(which Scheele had noticed) industrially. He declined to patent his 
process or make any profit from it, but instead turned it over to the 
French government. This action of Berthollet is characteristic of 
many scientists, who believe that because they were freely helped 
toward making their discoveries by the work of those who preceded 
them, so they should freely pass on the results of their own labors. 

Fig. 22. Chlorine bleaches indirectly. In the presence of sunlight 
chlorine reacts with water, liberating nascent oxygen. 

chlorine water chlorine wat^r ir]^iWi^)t 

oxygen gas- 


Patents on chemical processes are sometimes turned over by their 
discoverers to foundations, which license them to manufacturers. 
These foundations receive fees or royalties for the use of the process. 
These funds may be used by the foundations in making further 
research possible. 

Chlorine bleaches paper and textiles. "Dephlogisticated marine 
acid air," as chlorine was called before Davy showed it to be an 
element, was a chemical curiosity in 1774. Today, it is an indus- 
trial necessity, and more than two million tons of it are produced in 
this country annually. 

The chief use of chlorine is in the bleaching, or decolorizing, of 
textiles, chiefly cotton and linen, and of wood pulp. It cannot be 
used for bleaching silk or wool because it destroys their fibers. Chlo- 
rine bleaches indirectly, by oxidation. Oxygen, liberated by the ac- 
tion of chlorine on water, combines chemically with certain coloring 
matters and dyes, which, upon oxidation, become colorless. This re- 
action is: 

??? +!? ~* 2HC1 + t 

Chlorine is used as a bleaching agent either in the free condition, 
in chlorine water, or in some unstable chlorine compounds, such as 
bleaching powder, CaOCl 2 , and calcium hypochlorite, Ca (CIO) 2 . 
Laundry bleach, used extensively today, is made by adding liquid 
chlorine to a very cold solution of sodium hydroxide. This product 
is known as Javelle water and contains sodium hypochlorite, NaCIO, 
a salt of hypochlorous acid, HC1O. 

2NaOH + C1 2 -* NaCIO + NaCl + H 2 O 

Sodium hypochlorite decomposes easily, liberating atomic oxygen. 

Dilute solutions of sodium hypochlorite, in strengths ranging from 
four to six percent, are marketed as "Clorox," "Rose-X," and a num- 
ber of other trademarks. Such household bleaches are used very 
widely in home laundering and in removing stains. In bleaching with 
chlorine, care must be taken not to expose the materials to the action 
of the chlorine for too long a time, since continued action will make 
the fibers very weak. After the bleaching agent has been used, the 
fabrics should be thoroughly rinsed in fresh water to remove all 
traces of the bleach. In industrial bleaching, after removal from the 
bleaching tank, an antichlor, such as sodium thiosulfate (Na 2 S 2 O.<) , 
commonly called hypo, is added to remove the excess chlorine, and 
then the material is thoroughly rinsed in running water. 


Chlorine saves lives. Although it consumes only about ten percent 
of normal production, a second very important use of chlorine is in 
the purification of water. When chlorine is added to water, atomic 
oxygen, liberated from the water by chlorine, reacts with the bacteria 
present and kills them by oxidation. Only three pounds of chlorine 
are used for each million gallons of water. In the United States more 
than 75 percent of all drinking water is treated with chlorine. As a 
result, the death rate from typhoid fever, a disease caused by typhoid 
bacilli which may be present in drinking water, has been cut down 
considerably. This treatment also kills algae and other low plant 
and animal life. For field use, in areas where pure water is not avail- 
able, explorers, scouts, and others carry tablets of a chlorine-produc- 
ing substance such as "halazone" which contains Ca (CIO) 2 . 

Chlorine is also used as an antiseptic and disinfectant. A substance 
that checks the growth of bacteria is called an antiseptic, while a 
germicide kills the bacteria outright. A substance that either kills or 
checks the growth of bacteria is called a disinfectant. Since much ill- 
ness is caused by poisons, or toxins, formed by bacteria, it is impera- 
tive that these bacteria be killed or at least prevented from mul- 

"Zonite," a trade preparation containing NaCIO, is used as a 
general household disinfectant. Because chlorine disinfectants if used 
improperly may destroy body tissue, such preparations should be used 
with great care. Several widely used insecticides are chlorine com- 
pounds (see page 540) . 

And chlorine destroys human life, too. Chlorine played a double 
role during World War I. While chlorine disinfectants and water 
chlorination were saving thousands of lives, free chlorine in the 
form of clouds of poisonous gas was choking out many other 
lives. Chlorine, and later phosgene, COCL, and mustard gas, 
(C1CH 2 CH 2 ) 2 S, caused terrible suffering in World War I, even 
though the use of poisonous gases had been "outlawed" by a confer- 
ence at The Hague in 1907. 

How to test for the presence of a chloride. Free chlorine may be 
recognized by its characteristic color and odor, but compounds of 
chlorine cannot be identified so easily. Because chlorides are presejat 
in so many common substances, a simple test for a chloride is de- 

All chlorides are soluble in water, with the exception of the 
chlorides of silver, mercury, and lead. Because silver chloride is in- 
soluble in water, chlorides are recognized by their reaction with a 
solution of silver nitrate. When a solution of silver nitrate is added 


silver nitrate 

Fig. 23. Testing an unknown solution for 
the presence of a chloride. What will hap- 
pen to the white precipitate on exposure 
to light if a chloride is present? 

to a solution of a chloride, a white insoluble substance, a precipitate, 
forms. On exposure to light, the color of this precipitate changes 
gradually to purple and then to black. 

AgNO 3 + NaCl - NaNO 3 + AgCl | 

This color change, when silver nitrate is added to an unknown 
solution, suggests that probably a chloride is present. But some 
substances that are not chlorides form similar compounds. Therefore, 
the formation of such a precipitate on the addition of silver nitrate 
to a solution is not an entirely reliable test for a chloride. 

A chemical test for a substance has one important requirement 
it must be specific, that is, no other substance will react to the test in 
the same way as the substance for which the test was designed. For- 
tunately, silver chloride is insoluble in dilute nitric acid; other sub- 
stances that might at first be mistaken for it are soluble in dilute 
nitric acid and disappear at once when this acid is added. The addi- 
tion of nitric acid, therefore, is the final step in this test, distinguish- 
ing the chloride from other compounds. 

An element is known by the company it keeps. The halogens (salt 
formers) are a group of elements that resemble one another chemi- 
cally, and whose physical properties differ from one another in 
regular gradation, as shown in Table 5. Such a group of elements 
is called a chemical family. The members of the halogen family are 
fluorine, chlorine, bromine, iodine, and astatine, a radioactive ele- 
ment. Table 5 shows the relationship among the members of the 
halogen family. 

Making bromine in the laboratory. Like chlorine, bromine is pre- 
pared by the oxidation of its hydrogen compound by manganese diox- 
ide. A mixture of potassium bromide, sulfuric acid, and manganese 
dioxide is heated in a test tube (see Fig. 24) . The H 2 SO 4 reacts 
with the KBr, forming HBr, which is then oxidized by the MnO 2 . 
Although at room temperature bromine is a liquid, it is liberated as 























Greenish - 




BROMINE, Br 2 1 







IODINE, I 2 1 



Purplish - 


Least soluble 
of halogens 


a brownish vapor at the temperature of the experiment. As this 
brownish vapor passes into the water, some of the bromine dissolves; 
the rest collects as a layer of bromine under the water. This method 
of collecting pure bromine by distillation is relatively safe. 

2KBr + 2H 2 SO 4 + MnO 2 -> K 2 SO 4 + MnSO 4 + 2H 2 O + Br 2 

Great care should be taken in working with bromine, because it is 
poisonous and attacks the skin, causing severe burns. Particular care 
should be taken to protect the eyes from bromine vapor. 

Taking bromine from the sea. Most bromine is extracted from the 
minute percentage (0.0065%) of bromides present in sea water. 
Free chlorine replaces the bromine of the bromides. Some of our 
bromine is also obtained from the bromides found in salt wells and 
salt lakes. The principal chemical reaction is: 

MgBr 2 + C1 2 - MgCl 2 + Br 2 

Bromine helps engine efficiency. Much of the bromine produced 
in the United States is used in the manufacture of "Ethyl fluid," 
an anti-knock mixture composed of ethylene dibromide and tetra- 
ethyl lead (TEL) . Large quantities of bromine are also used in 
making silver bromide, the light-sensitive chemical that forms the 
most important part of the coating of photographic films. Bromine 
and bromine compounds are used also in making tear-gas bombs. 
Bromine is used in appreciable quantities as an oxidizing agent in 
the manufacture of certain dyes and drugs. 

Potassium bromide or sodium bromide acts as a depressant on the 
central nervous system. Their action is followed by drowsiness and 
even sleep. Such chemicals are called sedatives. They are used in the 





CaF 2 
MgBr 2 

Nal0 3 
Mgl 2 







Forms HF 


H 2 + F 2 
-*2HF + 




Forms HCI 


H 2 + CI 2 
-*2HCI + 




Forms HBr 


Br 2 -t-H 2 O 
-* 2HBr + O 

Less than 



Forms HI 


I 2 + H 2 

- 2HI + 




treatment of insomnia and asthma and are frequently found in head- 
ache and sleeping powders. Heavy, continuous doses of bromides may 
have harmful effects on the body. Bromides should be used only on 
the advice of a physician. 

Methyl bromide, CH ;5 Br, is used widely in commercial and indus- 
trial fumigation to kill insects and other low forms of plant and ani- 
mal life. It is used for this purpose in boxcars, warehouses, and food 
processing and packaging plants. 

How iodine is prepared in the laboratory. Like chlorine and 
bromine, iodine is prepared by the oxidation of its hydrogen com- 
pound by manganese dioxide. A mixture of potassium iodide, sulfuric 
acid, and manganese dioxide is heated in a test tube as shown in 
Fig. 25. 

SKI + 2H 2 S0 4 + Mn0 2 - K 2 SO 4 + MnSO 4 + 2H 2 O -f I 2 

The violet vapor of iodine, which is produced at the temperature 
of the reaction, condenses, forming purplish-black crystals on the in- 
side of the test tube. This process of collecting iodine is called sub- 

Fig. 24. Preparation of bromine in 
the laboratory. Why is bromine not 
collected in the same way as chlo- 




























- 187C 


Forms HF 


H 2 + F 2 
-* 2HF + O 



CaF 2 









Forms HCI 


H 2 + CI 2 
-* 2HCI + O 




BROMINE, Br 2 1 








Forms HBr 


Br 2 + H 2 O 
-* 2HBr + O 

Less than 


MgBr 2 

IODINE, I 2 f 



Purplish - 


Least soluble 
of halogens 

200 C 


Forms HI 


I 2 + H 2 O 
-* 2HI + 



NalO 3 
Mgl 2 

a brownish vapor at the temperature of the experiment. As this 
brownish vapor passes into the water, some of the bromine dissolves; 
the rest collects as a layer of bromine under the water. This method 
of collecting pure bromine by distillation is relatively safe. 

2KBr + 2H 2 SO 4 + MnO 2 - K 2 SO 4 + MnSO 4 + 2H 2 O + Br 2 

Great care should be taken in working with bromine, because it is 
poisonous and attacks the skin, causing severe burns. Particular care 
should be taken to protect the eyes from bromine vapor. 

Taking bromine from the sea. Most bromine is extracted from the 
minute percentage (0.0065%) of bromides present in sea water. 
Free chlorine replaces the bromine of the bromides. Some of our 
bromine is also obtained from the bromides found in salt wells and 
salt lakes. The principal chemical reaction is: 

MgBr 2 + C1 2 > MgCl 2 + Br 2 

Bromine helps engine efficiency. Much of the bromine produced 
in the United States is used in the manufacture of "Ethyl fluid," 
an anti-knock mixture composed of ethylene dibromide and tetra- 
ethyl lead (TEL) . Large quantities of bromine are also used in 
making silver bromide, the light-sensitive chemical that forms the 
most important part of the coating of photographic films. Bromine 
and bromine compounds are used also in making tear-gas bombs. 
Bromine is used in appreciable quantities as an oxidizing agent in 
the manufacture of certain dyes and drugs. 

Potassium bromide or sodium bromide acts as a depressant on the 
central nervous system. Their action is followed by drowsiness and 
even sleep. Such chemicals are called sedatives. They are used in the 



treatment of insomnia and asthma and are frequently found in head- 
ache and sleeping powders. Heavy, continuous doses of bromides may 
have harmful effects on the body. Bromides should be used only on 
the advice of a physician. 

Methyl bromide, CH,,Br, is used widely in commercial and indus- 
trial fumigation to kill insects and other low forms of plant and ani- 
mal life. It is used for this purpose in boxcars, warehouses, and food 
processing and packaging plants. 

How iodine is prepared in the laboratory. Like chlorine and 
bromine, iodine is prepared by the oxidation of its hydrogen com- 
pound by manganese dioxide. A mixture of potassium iodide, sulfuric 
acid, and manganese dioxide is heated in a test tube as shown in 
Fig. 25. 

2KI + 2H 2 SO 4 + MnO 2 -> K 2 SO 4 + MnSO 4 + 2H 2 O + I 2 

The violet vapor of iodine, which is produced at the temperature 
of the reaction, condenses, forming purplish-black crystals on the in- 
side of the test tube. This process of collecting iodine is called sub- 

KBr + MnO 2 
H 2 SO 4 

Fig. 24. Preparation of bromine in 
the laboratory. Why is bromine not 
collected in the same way as chlo- 


iHk" bromine 

Dow Chemical Company 

Methyl bromide used as a fumigant destroys all grain insects. 

limation. A substance is said to sublime when it passes directly from 
the solid state to the gaseous state and then condenses back to the 
solid state without passing through the liquid state. Camphor 1 is an- 
other substance that sublimes. 

The industrial preparation of iodine. About 90% of our iodine 
is obtained from brine that comes up with the oil in California oil 
fields. This brine contains Nal and MgL. Chlorine is passed through 
the brine and replaces the iodine. The iodine is recovered by adsorp- 
tion on activated carbon. (Adsorption is the clinging of molecules of 
one substance to the surface of another see page 327.) The princi- 
pal reaction is: 

MgI 2 + C1 2 -> MgCl 2 + I 2 

The rest of our iodine comes from NaIO ;5 , found as an impurity 
in Chile saltpeter, NaNO,. 

Iodine, too, saves lives. The chief use of iodine is in the prepara- 
tion of tincture of iodine, a two percent solution of iodine and po- 
tassium iodide in ethyl alcohol, which is an excellent antiseptic. As 
silver iodide, iodine is used to some extent in photography together 
with silver bromide. It is used also in the manufacture of iodoform, 
CHI :1 , a yellow powder used as an antiseptic, and in the manufacture 
of "Aristol," an improvement over iodoform. Iodine is also used in 
the production of certain dyes and methyl iodide, CHJ. 




Is iodine necessary to health? Iodine is an important constituent 
of the human body. There is a definite relation between the presence 
of iodine in the thyroid gland and the prevalence of certain disorders. 
The thyroid gland, located in the neck, secretes a compound called 
thyroxin, containing about 65 percent iodine, which helps to regulate 
the rate of oxidation in the body. 

When the thyroid gland receives too little iodine, goiter, an en- 
largement of the thyroid, results, caused apparently by the attempt of 
the gland to increase its size in order to produce more thyroxin. To 
offset this deficiency, iodides may be added to drinking water or 
about 0.02 percent of sodium iodide added to so-called iodized salt. 

Extreme underactivity ot the thyroid gland in newborn babies and 
young children may result in cretinism (kre'tm-izm) misshapen 
dwarfishness, low mentality, sluggishness, dullness, slow heart action. 
Synthetic thyroxin is used in the treatment of this thyroid disorder. 
Overactivity of the thyroid gland often produces the opposite effect 
the thin, nervous, highly energetic person, whose movements are 
quick, and whose heart action is rapid. See page 34 for a discussion 
of basal metabolism tests in diagnosis of thyroid disturbances. 

Iodine is also necessary to other forms of animal life. Large quan- 
tities of iodides are added to commercial feeds for chickens, cattle, 
dogs, cats, and other animals, and to fertilizers for forage crops. 

Replacement power of the halogens. If free chlorine is added to a 
solution of a bromide or an iodide, free bromine or free iodine is 
liberated. Free chlorine replaces the two less active halogens. 

2KBr + C1 2 -> 2KC1 + Br 2 
2NaI + C1 2 -> 2NaCl + I 2 

The addition of free bromine to a solution of an iodide releases 
free iodine. 

2KI + Br 2 -* 2KBr + I 2 

test tube containing I 

' H 2 SO 4 


Fig. 25. Laboratory preparation of 

iodine crystals 


However, the addition of iodine to a solution of either a bromide or 
a chloride has no effect, for the less active halogen will not replace 
the more active halogen from its compound. As mentioned earlier, 
the commercial preparation of bromine depends on the replacement 
power of chlorine. 

How we test for the presence of bromides and iodides. Many uses 
are made of the replacement power of the halogens. The tests for 
bromides and iodides are based on it. Chlorine water is added to a 
solution of the unknown salt, and a few drops of carbon disulfide, 
CS 2 , which is not soluble in water, are also added. If, after shaking 
the mixture, the carbon disulfide settles out as a distinct layer with 
a brownish-red coloration, then the original salt was a bromide, the 
free bromine liberated coloring the carbon disulfide. If the carbon 
disulfide acquires a purple coloration, the original salt was an iodide. 

Carbon disulfide is used because free bromine and iodine are much 
more soluble in this liquid than in water. Hence, most of the liber- 
ated bromine and iodine dissolve in the carbon disulfide, thus color- 
ing it much more than they would color water. Carbon disulfide, be- 
cause it is a better solvent, will extract any bromine or iodine from 
the water solution. This process of separation, frequently used in 
industry, is called extraction by partition. 

Fluorine, the most active of all the elements. Fluorine was not 
isolated until 1886. Because of its extreme chemical activity, which 
causes it to unite violently with metals, glass, porcelain, and water, 
its separation as a pure element was a very difficult undertaking. 
Finally, Henri Moissan succeeded by liquefying pure hydrogen 
fluoride, adding some potassium fluoride, and at a temperature of 
23C. passing an electric current through the mixture. Fluorine 
was liberated at the anode. 

The anode used industrially today is made of graphite, which is 
not attacked by this pale yellow gas. Fluorine is employed in making 
uranium hexafluoride used in atomic energy plants where many of 
the lubricants are chemically inert fluorocarbon compounds. The 
plastic, "Teflon," is another fluorine compound. The new rat and 
ground squirrel poison, 1080, is a fluorine compound, and sodium 
fluoride is used in some insecticides. 

Fluorine prevents tooth decay. The amount of tooth decay, or 
dental caries, has been found to vary directly with the quantities of 
fluorides present in the- local water supply. Too much fluoride pro- 
duces very hard but mottled teeth, a condition in which the enamel 
becomes discolored badly. When too little fluoride is present, there 
is much tooth decay. Water fluoridation (about one part NaF per 



million parts water) is now widely practiced to protect children up 
to about the age of 12 while tooth enamel is being formed. 

Fluorine refrigerants. A refrigerant is a substance used to absorb 
heat by changing from a liquid to a gas. In refrigeration, the material 
from which heat is absorbed is cooled. Almost as long as mechanical 
refrigerators have existed, their manufacturers have searched for bet- 
ter refrigerants. Something was wrong with nearly all the original 
refrigerants. They were either toxic, flammable, corrosive, or pos- 
sessed disagreeable odors. And then a family of compounds was de- 
veloped, and introduced in 1931 under the trademark "Freon." 
These compounds, produced by the halogenation of simple com- 
pounds of carbon and hydrogen, are far superior to sulfur dioxide, 
ammonia, ethyl chloride, and methyl chloride as refrigerants. All are 
practically odorless, nontoxic, nonflammable, and noncorrosive. 

The first of the "Freons" to be produced was dichlorodifluoro- 
methane, CC1,F,, and a later one was dichloroteArafluoroethane, 
C 2 C1 2 F 4 . There are several others, each with slightly different proper- 
ties that make it particularly well adapted to a special use. Because of 
their extreme volatility and the speed with which they penetrate 
every nook and cranny of a confined space, the "Freons" are used as 
propellants in dispersing insecticides. Aerosol bombs used for killing- 
household pests usually contain an insecticide and a liquefied "Freon" 
gas under pressure. When the pressure is released, the expanding gas 
quickly distributes the insecticide throughout a room or closet. 

This electronic device, called a "sniffer/ 1 is used to detect 
breaks in telephone cables through which moisture might 
enter. The cable to be tested is filled with a "Freon 1 ' gas. 
Then the sniffer is pulled along the cable from the ground. 
Escaping gas activates a "FreorT'-sensitive detector, ringing 
a bell. 

Bdl Telephone laboratories 

|P^'^-^ : ^^^^ 



Consumer Reports (Consumers Union Reports) , April, May, 
June, 1939. Consumers Union of U.S., New York. Excellent 
reports on various antiseptics. 

Harrow, Benjamin. Eminent Chemists of Our Time (2nd 
ed.) . D. Van Nostrand Co., New York, 1927. Read the life of 
Moissan and his isolation of fluorine. 

Sanders, Gardiner, and Wood. "Chlorine and Caustic Manu- 
facture/' Industrial and Engineering Chemistry, September, 
1953, pp. 1824-1835. Includes history, production figures, 
photos, and diagrams showing how diaphragm and mercury 
cells work. 


1. The halogens are a group of elements that resemble one 
another chemically, and whose physical properties differ in 
regular gradation. The halogen group is one of several such 
groups of elements. 

2. A substance that passes directly from the solid to the 
gaseous state and directly from the gaseous to the solid state 
is said to sublime. 

3. The ability of one element to replace another in that ele- 
ment's compounds is widely used in chemical reactions. The 
halogens may be listed according to their replacement power. 

4. The separation of a substance from a solution containing 
that substance by means of a better solvent is known as extrac- 
tion by partition. 

5. Like Berthollet, many scientists believe they should freely 
pass on to others the results of their own labors, and so help 
all humanity. 


Group A 

1. Who discovered C1 2 , and in what year was it isolated? 

2. Make a labeled diagram showing the laboratory prepara- 
tion of C1 2 . 

3. What is the function of MnO 2 in the laboratory prep- 
aration of C1 2 ? 

4. What other substance might be used instead of MnO 2 in 
the preparation of C1 2 ? 


<5. (a) .Give two reasons for collecting C1 2 by the displace- 
ment of air. (b) Why should any excess of C1 2 be passed into 

6. Write a balanced equation representing the laboratory 
preparation of Cl a . 

t . . . 

7. State four physical properties of C1 2 . 

8. Faraday liquefied C1 2 in 1823. It was the first gas to be 
liquefied. Suggest a reason. 

9. If a brightly burning paraffin taper is inserted in a bottle 
of C1 2 , a heavy black smoke is given off. Explain. 

10. By what process does C1 2 bleach? Explain. 

11. (a) Why is it necessary to rinse materials that have 
been bleached with C1 2 ? (b) Why cannot C1 2 be used to bleach 
silk and wool? 

12. Because of what chemical property are bleaching pow- 
der, CaOCl 2 , and Clorox, containing NaOCl, able to bleach? 

13. Find the percentage of C1 2 in bleaching powder. 

14. (a) Distinguish between the terms antiseptic and germi- 
cide, (b) What term embraces both? 

15. Make a list of all the uses of C1 2 that you know. 

16. Chloride of lime, or bleaching powder, has the formula 
CaOCl 2 . It is made from Ca (OH) 2 and C1 2 . Write the equa- 
tion for its preparation. 

17. Give a brief account of the preparation of some laundry 

18. Chlorine played a double role in wartime. Explain. 

19. How much CI 2 can be prepared by the action of 348 g. 
of MnO 2 on sufficient concentrated HC1? 

. . t . . . 

" 1 

20. (a) Describe fully the test for a chloride, (b) Write the 
equation for the reaction that takes place. 

21. Illustrate the statement, "An element is known by the 
company it keeps." 

22. (a) Compare the physical properties of the members of 
the halogen family, (b) Compare their chemical properties. 

23. In what ways do the halogens (a) resemble one an- 
other, and (b) differ from one another? 

24. List the halogens in the order of their chemical activity. 


25. Illustrate what is meant by the replacement power of 
the halogens. 

26. Does the following equation represent an actual chem- 
ical reaction? Explain. 2KC1 + Br 2 >2KBr + C1 2 

27. Describe two uses of the replacement power of the 

28. Make a labeled diagram of the laboratory preparation 

of Br 2 . 

29. (a) What is a sedative? (b) State three other uses of 


30. Describe the commercial preparation of Br 2 . 

Br r 

31. Using a labeled diagram, describe the laboratory prep- 
aration of I 2 . 

32. (a) What is sublimation? (b) What is tincture of 

33. How is iodine obtained for industrial use? 

34. How is the most active of all the chemical elements 

35. (a) What is one cause of mottled teeth? (b) What is 

36. What is the relationship of F 2 to dental caries? 

37. (a) What compounds of F 2 are superior refrigerants? 
(b) Why? 

Group B 

38. Compare C1 2 and O 2 with respect to: (a) chemical 
activity, (b) behavior with H 2 , (c) valence. Explain each 
answer fully. 

39. (a) What is the relationship between lack of iodine 
and goiter? (b) Why is goiter not so prevalent in New York 
City as it is in some other parts of the United States? 

40. A bottle of tincture of iodine was found, after long 
use, to contain only a dark solid, (a) Would it be safe to use 
it after adding pure ethyl alcohol? (b) Explain your answer. 

41. If you had some Nal, how would you prepare a solution 
of tincture of iodine from it? 

42. I 2 is produced from California oil-well brines and from 
other brines in Michigan. Can you suggest a way in which 
this I 9 is extracted? 


43. How much fluorine would be needed to make one ton 
of dichlorotetrafluorethane, C 2 C1 2 F 4 ? 

44. (a) What factors must a manufacturer consider when 
he chooses raw materials for use in manufacturing a substance 
on a large scale? (b) Show how these factors apply in the man- 
ufacture of C1 2 . 


1. Make a study of advertisements of antiseptics and dis- 
infectants in magazines and newspapers. What appeals are 
made in these advertisements to induce you to buy a particular 
brand? Are the appeals chiefly scientific, pseudo-scientific, fear- 
provoking, or do they appeal chiefly to your pride, sense of 
superiority, your desire for social approval, and so forth? Il- 
lustrate your report with actual advertising copy. 

2. Ink eradicators frequently contain two solutions. No. 1 
contains a solution of a very weak acid and No. 2 contains a 
solution of sodium hypochlorite. Prepare such an ink eradi- 
cator using vinegar or citric acid for No. 1 and Javelle water 
or Clorox for No. 2. Demonstrate its use before your class. 

3. Victor Meyer, an eminent German chemist, prepared a 
compound in 1886 which is now known as mustard gas. Meyer 
was the first chemist to prepare this chemical during his the- 
oretical investigations. Thirty years later chemists, cooperating 
with the Germany military machine, made this compound 
available for use as a poison gas. Should society take a hand in 
suppressing such discoveries which might be used against man- 
kind? Write a brief paper either in favor of this point of view 
or against it, or arrange in class for a discussion or debate on 
this topic. 

4. Study the quotation at the beginning of this chapter. 
Prepare a brief report to the class on the meaning of this 
quotation. Illustrate your report with an example from the 
history of science. 



. . . Our experiences and observa- 
tions alone never lead to finalities. 
Theory, however, creates reliable 
roads over which we may pursue our 
journeys through the world of ob- 
servation. Anton Reiser, 1930 

The electron theory gives us a clearer picture of matter and its 
changes. Formulating accurate theories takes remarkably clear in- 
sight, courage, and creative imagination. The theories and principles 
of science are among the most truly creative products of the mind of 
man. The atomic theory of Dalton is one of the great theories upon 
which modern chemistry is built. It shows us that atoms do not com- 
bine in a haphazard, irregular manner but form molecules in ac- 
cordance with unvarying natural laws. However, it gives us no idea 
why this is so. 

A more recently developed theory, which is called the electron the- 
ory, supplements the atomic theory and gives this explanation. It pro- 
vides answers also to such questions as why the extreme chemical 
activity of fluorine, the comparative inactivity of nitrogen, and the 
inertness of argon; why elements and radicals possess the valences 
that they have, and many other questions. The electron theory is the 
fruit of many scientists who worked in many countries. It made possi- 
ble the atomic age. 

Static electricity. About 2600 years ago, the Greeks discovered that 
when amber is rubbed with cloth, it becomes capable of attracting 
tiny bits of straw or dry leaves. Through the centuries, men have 



rubber rod 

Fig. 26. Demonstration of static electricity. 

glass rod 

discovered that other materials may be given this property. Glass, 
when rubbed with silk, or hard rubber, when rubbed with fur, will 
also attract light objects. The force which causes this attraction was 
named electricity from the Greek word for amber. Today, we refer 
to the electricity caused by friction or rubbing as static electricity. 

Benjamin Franklin attempts to explain negative and positive 
electricity. Suspend a pith ball by a silk thread. Touch it with a hard- 
rubber rod which has been rubbed with fur. As soon as contact is 
made, the pith ball will be driven away, or repelled, from the rod. 
Then bring near the pith ball a glass rod which has been rubbed 
with silk. The ball will be attracted toward the glass rod. 

This simple experiment demonstrates that the glass and the hard- 
rubber rods were charged with opposite kinds of electricity. When 
the neutral pith ball was touched by the hard-rubber rod, it became 
charged with the same kind of electricity as the rod. It was then 
pushed away from the rod, proving that objects with the same electric 
charges repel one another. However, the ball was drawn toward the 
glass rod, proving that objects with opposite electric charges attract 
one another. 

In 1747, Benjamin Franklin, one of the most versatile men Amer- 
ica has ever produced, received a static electricity machine Irom a 
friend in England. Franklin, son of a soap-maker who had fled from 
England because of religious persecution, was then 41 years old. Be- 
cause of a very successful business career, he was rich enough to re- 
tire from business and devote himself to scientific experimentation. 
He had already organized the first scientific society in the New 
World, an organization which later became the American Philosophi- 
cal Society. 

Franklin performed many experiments with his static electricity 
machine, and that same year he announced his own views on the 
nature of electricity. He wrote: "The electric fire (electricity) is not 
created by friction, but collected, being really an element diffused 
among matter. The electrical matter consists of particles extremely 
subtile. . . . Hence have arisen some new terms among us: we say 
B is electrised positively; A, negatively. Or rather B is electrised plus; 



A, minus/' Franklin was the first person to use these present-day 
terms in referring to electricity. He later induced his good friend 
Priestley to write a history of electricity, and thus, in part, directed 
Priestley's scientific career. 

Franklin's electric theory was not altogether correct, for he be- 
lieved that if a body has too much electricity it is charged positively 
(-f ) ; if it has not enough, it is charged negatively ( ); and if it has 
just enough, it is neutral. Even though his ideas were not altogether 
correct, his reasoning and his terminology for electricity were more 
modern than those of any other eighteenth-century scientist. So 
great was the creative imagination of Franklin that he came very 
close to arriving at the modern concept of the electric nature of mat- 
ter, a point of view reached only after some 150 years of further ex- 

Today we refer to the charge which was produced on the hard- 
rubber rod mentioned above as negative ( ) , and that produced on 
the glass rod as positive (+) . 

The electron is discovered. Almost a century and a half after 
Franklin, William Crookes, an Englishman, studied the effect of 
passing a current of high-voltage electricity through a glass tube from 
which nearly all the air had been pumped. He noticed that a beam of 
light issued from the negative plate, or cathode, of the tube. When a 
magnet was brought near the tube, the beam would bend. Since ordi- 
nary light is not affected by a magnet, the beam showed a property of 
matter rather than of light. 

Another Englishman, J. J. Thomson, undertook to explain the 
strange behavior of these cathode rays. In 1897, after 20 years of bril- 
liant research, he announced his results. He said that cathode rays 
are composed of particles of negative electricity, torn away from the 
atoms of the air in the tube. To these particles, Thomson gave the 
name electrons. The cathode ray was bent because the negative elec- 
trons would be attracted by the positive pole and repelled by the 
negative pole of a magnet. 

Fig. 27. Crookes' tube. Notice how the positive pole 
of the magnet deflects the cathode ray. 

stream of electrons 

Crookes' tube 

Joseph John Thomson was 
born at Manchester in 1856. 
He followed Rayleigh as head 
of the Cavendish Laboratory 
of Experimental Physics at 
Cambridge University. 

Brumt font t t c, 

The discovery of the electron also explained the phenomenon of 
static electricity. When one object is rubbed with another, electrons 
are transferred. Thus, when glass is rubbed with silk, the glass loses 
electrons to the silk and is left with a positive charge; amber, when 
rubbed with fur, takes electrons from the fur and becomes negatively 

Thomson's discovery of the electron completely upset the theory 
that the atom is the smallest unit of matter, since electrons were 
found in all atoms. It also proved that the idea of an indivisible atom 
was inaccurate. 

Subsequently, Robert A. Millikan, an eminent American scientist, 
succeeded in computing the mass of a single electron. He found it 
to be about y^Vr * ^ ie mass * one hydrogen atom. (Mass is the 
amount of matter that a substance contains. It does not vary, as does 

lirown Brothers 

Robert Andrews Millikan (1868- 
1 953) won the Nobel physics prize 
in 1923 for his work in isolating 
and weighing the electron. 




electron (outside the nucleus) 
proton (in the nucleus) 

Fig. 28. The hydrogen atom 

weight, with the gravitational pull of the earth. However, weight is 
dependent upon the amount of matter a substance contains.) 

The proton is discovered. After Thomson discovered the electron, 
one of his students, Ernest Rutherford, began to ponder over the na- 
ture of the rest of the atom. An atom itself is electrically neutral. 
Surely, he thought, in the electrically neutral atom there must be 
some positive electricity to counterbalance the negative electron. 

After much research, which began in 1911, Rutherford determined 
that all atoms contain one or more particles of positive electricity 
which he named the proton. The proton, which is a positively 
charged atom of hydrogen, is 1837 times as heavy as the electron. 

The structure of the atom according to the electron theory. Atoms 
of the various elements contain different numbers of electrons and 
protons. In general, however, the arrangement of these particles fol- 
lows a similar pattern in all elements. Rutherford gave us the first 
picture of the structure of the atom. It resembles our own solar sys- 
tem with its sun and revolving planets. The "sun" of the atom is 
called the nucleus. It is composed partly of protons around which, 
at a relatively great distance, revolve planetary electrons. A diagram 
of the simplest of the atoms, that of hydrogen, is shown above. 

It was Rutherford who found that electrons and protons were not 
evenly distributed in the atom but that the heavy protons were all 
located in the center. He shot helium atoms (alpha particles) 
through a cluster of nitrogen atoms and photographed the results 
by means of a cloud chamber and special camera devised by 
C. T. R. Wilson, another English scientist. He found that only in 
an extremely few cases the path or fog track of the helium bullet 
was not straight, but was thrown sharply off its course. On the basis 
of the volume of the nitrogen atom and the ratio of straight fog 
tracks to bent ones, he calculated that the protons must all be con- 
Fig. 29. (left) Apparatus for producing the cloud effect. 
The upper chamber is filled with nitrogen, (right) The 
fog track of an alpha particle. 

.water vapor condenses 


track of particle 
suddenly released 


water. . . 

is compressed . 


centrated in the tiny nucleus of the atom. The diameter of the nu- 
cleus is about 10 o*ooo that of the whole atom. Helium bullets pass- 
ing through the rest of the atom met no solid, positively charged 

Size of the atom. Most of us are baffled in attempting to imagine 
the size of the particles within the atom and the distances between 
them. The diameter of a hydrogen atom is about 40,000,000 inch. 
If an atom were magnified about 30 trillion times, its diameter would 
be about 10 miles. At the center would be the nucleus, about the size 
of a tennis ball. The electrons, each about the size of a hazelnut, 
would revolve about the nucleus in orbits in somewhat the same way 
as the planets of our solar system revolve about the sun. 

This description is actually an oversimplification of the structure 
of the atom. However, it tells us an important fact. Atoms, which 
compose every element and compound, are largely empty space! 

The neutron: A notable scientific prediction. All of the protons 
of an atom are located in the nucleus. However, not all of the elec- 
trons of an atom are planetary; some electrons, too, are found within 
the nucleus. But how can negatively charged electrons and positively 
charged protons exist side by side in the nucleus? To explain this, 
Dr. W. D. Harkins predicted the existence of a new particle. An 
electron within the nucleus, he said, does not exist as a separate par- 
ticle, but is combined with a proton forming an electrically neutral 
particle which he named the neutron. Since the weight of the elec- 
tron is extremely slight, it may be disregarded in figuring the weight 
of the neutron. The neutron has been found to have about the same 
mass as the proton. 

In 1920, the existence of the neutron was theoretically established 
by W. D. Harkins of the University of Chicago. Twelve years later, 
it was actually discovered by James Chadwick, working in Ruther- 
ford's laboratory. The ultimate verification of Harkins' work by 
Chadwick demonstrates the value of pure theory and creative imag- 
ination in science. For his part in the discovery, Chadwick was 
knighted in 1944. 

The electron and proton are so close together in the neutron that 
the volume of the neutron is millions of times smaller than that of 
any atom. It has, therefore, an extremely high density. This fact has 
been used to explain the extremely high densities of certain stars. 
Since the neutron is electrically neutral, electric forces do not repel 
it. Therefore, the neutron has great penetrating powers. 

We may say that, in general, all matter is composed of three kinds 
of fundamental electric particles: electrons, protons, and \neutrons. 


Scientists have discovered other electric particles within the atom, 
as shown in Table 8, but the three mentioned here are considered 
the most important. 

How the structure of an atom may be represented graphically. 
The atom of each element contains a particular number of electrons, 
protons, and neutrons different from the number in the atom of 
every other element. Two American scientists, Lewis and Langmuir, 
developed a theory of the arrangement of planetary electrons which 
explains why each element has different chemical properties. 

According to this theory, the electrons outside the nucleus (plane- 
tary electrons) arrange themselves in successive rings, or shells. The 
first ring is complete when it contains 2 electrons; the second ring 
is complete when it contains 8 electrons; the third ring is complete 
when it contains 18 electrons; the fourth, when it contains either 18 
or 32 electrons. However, the outermost ring never contains more 
than eight electrons. According to Lewis and Langmuir, an atom 
with 30 planetary electrons will have its first ring complete with 
2 electrons, second ring complete with 8 electrons, third ring com- 
plete with 18 electrons, and fourth ring incomplete with 2 electrons. 
What is the electron structure of an atom with 48 planetary electrons? 

The number of protons in the nucleus of the atom is equal to the 
number of planetary electrons. This equality keeps the atom elec- 
trically neutral. Thus, an atom with 30 planetary electrons will have 
30 protons in its nucleus. 

Of all the elements, hydrogen has the simplest atom. Its nucleus 
consists of one proton. Revolving about this nucleus is one planetary 
electron. The nucleus of the helium atom contains two protons (2+) 
and two neutrons (2n) ; two planetary electrons (2) revolve about 
it. Fig. 30 shows how we may graphically represent the helium atom 
and also the chlorine atom (atomic number 17) . 

The periodic table of Mendeleyeff. In 1869, Mendeleyeff (men'- 
de-la'ef) , a Russian chemist, published a table of the elements ar- 
ranged in order according to their increasing atomic weights. He 
noticed, that when arranged in this manner, the elements fell into 

Fig. 30. Diagrams of the helium and chlorine atoms. 

Helium Chlorine 





































eight distinct groups. Within each group, the elements have similar 
physical and chemical properties. 

Let us examine part of this table carefully. Note that hydrogen is 
omitted, and that the table begins with helium, the element with the 
next heaviest atomic weight. Lithium (Li) , with an atomic weight 
of 6.940 follows, and so on through fluorine (19.0) . These eight dis- 
similar elements comprise one series or period. 

The element following fluorine in order of atomic weight is neon 
(Ne) . It has chemical properties similar to those of helium (He) and 
falls directly below it in the table. Directly below helium is argon 
(A) with similar properties. Thus helium, neon, argon, and certain 
succeeding elements comprise Group Zero. This group of elements 
is known as the family of inert gases (see Chapter 7) . The elements 
in Group VII are known as the halogen family (see Chapter 10) . 
The other groups are also made up of closely related elements. 

Moseley discovers the law of atomic numbers. Prior to 1912, the 
numerical position of an element in a table of atomic weight was 
called the atomic number of the element. It occurred to Rutherford 
that this number might also represent the number of protons in the 
nucleus of the atom. One of his students, Henry G. J. Moseley, un- 
dertook to find out whether Rutherford's idea was valid. 

Moseley's experiments bore out Rutherford's theory. The atom 
of each element was found to contain a number of protons in its nu- 
cleus corresponding to the element's numerical position in the peri- 
odic table of atomic weights. The hydrogen atom, which appears first 
in the periodic table, has only one proton in the nucleus of its atom; 
uranium, which appears in the ninety-second position in the periodic 
table, has 92 protons in the nucleus of each atom. Moseley showed 
that the atomic number of any element is equal to the number of free 
protons in the nucleus of its atom. This is known as the law of atomic 
numbers. Since the number of free protons in the nucleus is equal 
to the number of electrons around the nucleus, the atomic number 
of an element is also equal to the number of planetary electrons. 

For the first 17 elements in the table of atomic numbers, it is help- 
ful to remember that the atomic number is equal to half the atomic 






weight (disregarding fractions). Thus the atomic number of chlo- 
rine (atomic weight 35.457) is 17. 

The periodic table of atomic numbers. Mendeleyeff's periodic 
table based upon atomic weights served science for 50 years. In 1912, 
however, it was displaced by the new periodic table developed by 
Moseley from his law of atomic numbers. 

Moseley's table is more fundamental than MendeleyefFs and easily 
accounts for some of the discrepancies in the latter. For example, the 
element argon has an atomic weight of 39.944 and the element potas- 
sium an atomic weight of 39.10. Argon should, therefore, follow 
potassium in the table based on atomic weights. But the properties 
of argon put it in the group of inert elements, preceding potassium. 
Moseley's researches, which showed that the atomic number of argon 
is 18 and that of potassium is 19, eliminated this problem. When 
Moseley developed his table, all of the elements had not been discov- 
ered. Therefore the atomic numbers did not run entirely consecu- 
tively from one to 92. Since then the missing elements have been dis- 
covered, as you may see from the modern periodic table below. 

A new definition of atomic weight. As we have learned, the elec- 
tron, for most purposes, may be considered weightless, and the proton 
and neutron may be considered equal in mass. Thus, the most abun- 
dant hydrogen atom, which contains a single proton and no neutrons, 
may be considered equal to one proton in weight; the helium atom, 
which contains two protons and two neutrons, weighs four times as 
much as the hydrogen atom. We may say that the atomic iveight of 
hydrogen is one and the atomic weight of helium is four. Thus, ac- 
cording to the electron theory, the atomic weight for each element 
may be defined as the sum of the protons and neutrons in the nucleus 
of an atom of that element. Look at the diagrams of various atoms in 
Fig. 32. What is the atomic number of each? the atomic weight? 

How the electron theory explains isotopes. In 1815, William 
Prout, a London physician, announced the theory that all the chem- 
ical elements are made up of groups of hydrogen atoms only. Prout's 
theory was not taken seriously for 100 years until Moseley's work 
on atomic numbers made Prout's conclusion more plausible. 

Since the nucleus of an atom of any element is composed of only 


'Following lanthanum are 1 4 elements known as the rare earth elements (at. no. 58-71). 
* * Following uranium are eight newly created elements of the actinide series (at. no. 93-100). 









52.01 34 



50.95 33 









95.95 52 






180.88 83 


183.92 84 


186.31 85 








Henry 6. J. Moseley (1887- 
1915), a pupil of Ernest 
Rutherford, discovered the 
law of atomic numbers in 
1912. His brilliant career was 
ended by his death at 27 
during World War I. 

neutrons and protons, and since the weight of each of these units 
is really the weight of the hydrogen atom, it may have occurred to 
you that the atomic weights of all the elements ought to be whole 
numbers. But the fact that many atomic weights, for example, chlo- 
rine (35.457) , are not whole numbers could riot be brought into 
harmony with this idea. 

In 1913, Theodore W. Richards found two different kinds of lead 
with atomic weights of 206 and 207, respectively, and, in the same 
year, two kinds of neon with different atomic weights were reported 
also. The name isotopes was given to atoms of the same element hav- 
ing the same chemical properties but different atomic weights. Dis- 
coveries of isotopes of many other elements soon followed, one of 
which (tin) is now known to have as many as 10 stable isotopes and, 
hence, 10 different atomic weights. 

The discovery of isotopes removed the obstacles to the acceptance 
of Prout's idea. For elements, as we know them, are really mixtures 
of isotopes having different atomic weights, each of which is a whole 
number. Thus ordinary chlorine gas is really made up of some atoms 
with an atomic weight of 35, other atoms of atomic weight 37, and 
still others of atomic weight 39. Its accepted atomic weight, 35.457, 
is the average of the atomic weights of the three different weights of 
chlorine atoms in any sample of the gas. Here was another startling 
discovery which helped to destroy Dalton's idea of an atom whose 
atomic weight never changed. 

The electron theory explains isotopes as caused by a different num- 
ber of neutrons in each kind of atom. Thus, isotopes of chlorine with 
atomic weights of 35, 37, and 39 behave alike chemically because 
they have the same arrangement of planetary electrons. They differ 

Fig. 31. Diagrams of the isotopes of chlorine. 


35 ' 










in weight because of a difference in the number of neutrons in their 
nuclei. We must, then, redefine the term element. A substance of 
which all atoms have the same atomic number is an element. 

How the electron theory explains valence. From the diagram of the 
chlorine atom (Fig. 31) , you can see that its outermost ring contains 
7 electrons. One more electron is needed to make the 8 electrons 
needed to complete this ring. 

An atom whose outermost ring is nearly complete has a tendency 
to borrow enough electrons to complete this ring. An atom whose 
outermost ring has few electrons tends to lose electrons. The number 
of electrons gained or lost by an atom of an element is the valence of 
that element. Since the chlorine atom needs to borrow only 1 elec- 
tron to complete its outer ring, its valence is 1. In borrowing this 
electron it becomes negatively charged and, as a result, the valence 
of chlorine is negative. Hence, the valence of chlorine is 1. 

An atom that lends electrons becomes positively charged. Hence, 
elements whose atoms lend electrons have positive valences. Thus, 
the sodium atom with an atomic number of 11 (roughly half of 
22.997) may be pictured as shown in Fig. 32. As you see, the outer- 
most ring contains 1 electron which the atom may lend. Hence the 
valence of sodium is -f-1. 

An atom whose outer ring is complete will neither lend nor bor- 
row electrons. Elements whose atoms are of this type have a valence 
of 0. The atom of neon is shown in Fig. 32. 

Flectrons in the outermost ring of an atom, which may be either 
borrowed or lent, are called valence electrons. 

How the electron theory explains metals and nonmetals. A metal 
is a lender of electrons. That is, the outermost ring of an atom of a 
metal has less than four, or half the number (eight) required to 
complete it. When such an atom lends electrons, it necessarily be- 
comes positively charged. The valence of metals, therefore, is con- 
sidered positive. 

The atom of a nonmetal is a borrower of electrons. That is, its 
outermost ring has more than four, or half the number of electrons 
(eight) required to complete it. By borrowing electrons, such an 
atom becomes negatively charged. The valence of nonmetals is, 
therefore, negative. 

Fig. 32. Diagrams of various atoms. 







If the outermost ring of an atom of an element has just half the 
number of electrons required to complete it, it may either borrow or 
lend electrons. Such an element is said to be amphoteric. A common 
example of an amphoteric element is carbon (atomic weight 12) , 
whose diagram appears in Fig. 32. 

How the electron theory explains electric currents. According to 
'modern theory, an electric current is a flow of electrons. Atoms of 
metals, such as copper, silver, and gold, are good conductors of elec- 
tricity, because some of their electrons are held loosely, and can move 
freely through the solid. In general, nonmetals are poor conductors, 
because their electrons are not held as loosely as those of metals. 

How the electron theory explains chemical activity. An atom 
tends to complete its outer ring of electrons. If an element such as 
neon or argon already has its outer ring complete, that element is 
inert. That is, its atom will not lend or borrow electrons, and hence 
the element is completely inactive chemically. In general, the smaller 
the number of electrons an atom must either borrow or lend to com- 
plete its outer ring of electrons., the greater is the chemical activity 
of that atom. 

An atom, then, with either 1 or 7 electrons in its outermost ring 
is extremely active. Such an atom is fluorine, whose atomic weight 
is 19, and atomic number is 9. It has 7 electrons in its second ring 
and will borrow 1 more electron. Potassium, whose atomic weight 
is 39, has only 1 electron in its fourth ring, and hence can lend only 
1 electron. Atoms such as those ot oxygen and magnesium have 2 
electrons to borrow or lend and are quite active. Atoms of nitrogen 
and aluminum have 3 electrons to borrow or lend and are not very 
active. Generally, the farther away the outer ring of an atom is from 
the nucleus, the less is the attraction of the nucleus for its electrons. 

This helps explain the chemical behavior ot metals. For exam- 
ple, potassium is more active than sodium since the electron in its 
fourth ring can be lost more easily than the electron in the third 
ring of sodium. Conversely, the closer the outer ring of an atom is 
to the nucleus, the stronger is the attraction of the nucleus for its 
electrons. This fact helps explain the chemical behavior of non- 
metals. Fluorine is more active than iodine since its second ring has 

Fig. 33. Chemical union of sodium and chlo- 
rine according to the electron theory. 



sodium chloride 


a greater attraction for the electrons of other atoms than does the fifth 
ring of iodine. Both of these rules should be considered rough guides. 
Many exceptions occur since the whole problem of chemical activity 
is quite complex. 


1 . The at. wt. of sulfur is 32. Make a picture of its atom ac- 
cording to the electron theory, and explain its valence and 
chemical activity. 

2. With the aid of a diagram, show why helium, at. wt. 4, 
is inert. 

3. With the aid ot diagrams, show why lithium, sodium, 
and potassium belong to the same family ot elements. 

4. What are the valence and chemical activity of an element 
whose outer ring contains 4 electrons? 

5. With a diagram, explain the valence of the (OH) radical. 

How the electron theory explains chemical union and electro- 
valence. Chemical activity is the tendency ot atoms to complete their 
outer rings and form stable compounds. Chemical union is, there- 
fore, the shifting or sharing of electrons in the outer electron rings 
until a stable condition is reached. 

A metallic atom with a valence of + 1 exhibits a strong attraction 
tor a non metal lie atom whose outer ring needs 1 electron to com- 
plete it. For example, the union of sodium, a very active metal with 
a valence ot +1, with chlorine, a very active nonmetal with a valence 
of 1, may be represented thus: 

The extra electron on the outer ring of the sodium atom shifts 
over to the vacant space in the almost complete outer ring of the 
chlorine atom. Now the two outer rings are both complete and the 
resulting compound, sodium chloride, is very stable. Its water solu- 
tion conducts electricity. Compounds formed by the shifting of single 
electrons are polar or ionic compounds. Their valence is called elec- 

Fig. 34. Formation of a molecule of fluorine 
according to the electron theory. 

shared pair of electrons 


How the electron theory explains covalent compounds. In many 
chemical reactions, there is no actual shifting of single electrons but 
rather an equal sharing of a pair or pairs of electrons. The com- 
pounds formed are nonpolar compounds and the valence is called 
covalence. This kind of combination is generally very strong and the 
molecules so formed hold together well. Nonpolar compounds are 
stable and generally do not conduct electricity. Many organic com- 
pounds, such as alcohol, glycerin, and sugar, are nonpolar. 

The fluorine molecule (see Fig. 34) illustrates the sharing of a 
pair of electrons. There is neither a gain nor a loss of electrons 
simply a sharing. Molecules of other gases that consist of 2 atoms, 
such as oxygen and chlorine, exhibit this same sharing of electrons. 
Polar and nonpolar compounds are discussed more fully on page 238. 

How the electron theory explains oxidation and reduction. In 
the equation CuO + H 2 -^ Cu + H 2 O, copper oxide is reduced to 
copper and hydrogen is oxidized to water. The valence of Cu has 
changed from plus two (in CuO) to zero in free copper, and the cop- 
per has gained 2 electrons (Cu++ + 2e - Cu) . The hydrogen has 
lost an electron and changed to H+ in H 2 +O (H e - H + ) . 

Oxidation has been defined thus far as the union of a substance 
with oxygen. Reduction, similarly, has been defined as the removal 
of oxygen from a compound. From the viewpoint pf the electron 
theory these terms take on much broader meanings. A loss of elec- 
trons resulting in an increase in the valence of an element is called 
oxidation. A gain of electrons resulting in a decrease in the valence 
of an element is called reduction. 

Oxidation and reduction in this broader sense need not involve 
either oxygen or hydrogen. Thus, in the replacement of the iodine 
in potassium iodide by chlorine, as explained in the preceding chap- 
ter, we have an oxidation-reduction reaction. According to the elec- 
tron theory, this reaction is: 

K+I- + Cl -> K+C1- + 1 

The iodine in KI has lost an electron and changed to free iodine, 1. 
It has been oxidized by chlorine which has gained an electron and 
changed from free chlorine to the negative chloride ion., Cl~ (see 
page 232) . 

How the electron theory aids in balancing equations. We may see 
from this oxidation-reduction equation how each atom of iodine lost 
an electron and was oxidized and how each free chlorine atom gained 
an electron, forming a chloride ion with a negative charge. This was 
a relatively simple reaction. But how is a more complex oxidation- 


reduction equation balanced and how does our understanding of the 
electron theory help us to find correct coefficients to use in balanc- 
ing the equation? Let us consider some examples: 

EXAMPLE A: It is desired to reduce ferric chloride, FeCl 3 , to 
ferrous chloride, Fed,, by the use of the evil-smelling gas, 
H 2 S a common reducing agent. In the reaction, free sulfur 
and hydrochloric acid are formed also. 

1) Write the unbalanced equation for the reaction. 

FeCl 8 + H 2 S - FeCl 2 + S + HC1 

2) Select the atoms which, according to the electron theory, 
are either reduced or oxidized, that is, either gain or lose elec- 
trons. Place the valences involved at the upper right of the 
symbol of each atom. 

Fe+++Cl 3 + H 2 S -> Fe++Cl 2 + S + HC1 

3) By inspection, we see that each atom of iron gains one 
electron and is reduced; each atom of sulfur loses two electrons 
and is oxidized. The number of electrons involved in the oxi- 
dation of H 2 S is twice as great as the number of electrons in- 
volved in the reduction of one molecule of FeCl 3 . Therefore, 
twice as many molecules of FeCl 3 must have been reduced. 
Indicate this by writing the coefficient 2 before FeCl 3 in the 
equation. Then balance the equation by the methods you have 
already learned. 

2FeCl 8 + H 2 S - 2FeCl 2 + S + 2HC1 

In actual practice, chemists usually consider only the changes in 
valence of the atoms involved rather than the shifting or sharing of 
electrons. The sum of the changes in valence on both sides of the 
arrow must be equal. Consequently, this method of balancing equa- 
tions is known as the valence-change method. 

EXAMPLE B: Using the valence-change method, write the equa- 
tion for the reaction between the oxidizing agent, potassium 
permanganate, and hydrochloric acid. The products are potas- 
sium chloride, manganese chloride, water, and free chlorine. 

1) Write the unbalanced equation, omitting the subscript 
from the free chlorine. It is not necessary to indicate the 
valence of those atoms which retain the same valence through- 
out the reaction. 

4 + HC1- - KC1 4- Mn++Cl 2 + H 2 O + Cl 


2) Note that the following changes occur: 

a) Each atom of manganese in KMnO 4 gains five elec- 
trons in combining with chlorine to form MnCl 2 ; therefore the 
valence loss of the manganese is five. 

(b) Each atom of chlorine loses one electron in becoming 
free chlorine; therefore the valence gain of each atom of lib- 
erated chlorine is one. 

3) The five valences lost by the manganese must be balanced 
by the valence gained by the chlorine since no other element 
changes valence in the reaction. To do this, we must show that 
for each molecule ot KMnO 4 used, five atoms of Cl are lib- 
erated. Therefore, write the coefficient 5 before the free Cl: 

KMn0 4 + HC1 -> KC1 + MnCl 2 + H 2 O + 5C1 

4) As we know, free chlorine exists as a molecule composed 
of two atoms, therefore we add the subscript, 2, to Cl. In order 
to keep the valence changes in balance, we consider that two 
molecules of KMnO 4 are used in liberating five molecules of 
C1 2 . Therefore, write the coefficient 2 before KMnO 4 . 

2KMn0 4 + HC1 -> KC1 + MnCl* + H 2 O + 5C1 2 f 

5) Complete the balancing of the equation in the usual 
manner. Check your work carefully. 

2KMnO 4 + 16HC1 -> 2KC1 + 2MnCl 2 + 8H 2 O + 5C1 2 T 

From this discussion, we can formulate the following method of 
balancing oxidation-reduction equations: 

1) Write the unbalanced equation. 

2) Write the changes in valence at the upper right of the 
symbols of the atoms that are oxidized and those that are re- 

3) Find the valence gains and losses and decide on the coef- 
ficients that will make them equal. Remember that molecules 
composed of two atoms may be involved and adjust the coeffi- 
cients accordingly. 

4) Complete the balancing of the equation by the usual 

Changing theories and the spirit of modern science. In 1924, the 
electron theory, backed by a mass of experimental evidence, was 
quite generally accepted. In that year, Prince de Broglie (de bro'y') 
suggested that the electron is not merely a particle of electricity, as 
the electron theory explained, but like light, is composed of, possesses, 
or perhaps is attended by, a group of waves. 


By 1927, two Americans, Davisson and Germer, proved de 
Broglie's theory experimentally, showing that both electrons and 
protons possess a property of light (a wave phenomenon) . In 1929, 
de Broglie was honored with the Nobel prize in physics for his theory. 

The theory that matter possesses wave properties created an up- 
heaval in the existing theories. However, we still find it convenient 
to regard the electron, proton, and neutron as tiny individual par- 
ticles. If new facts show we are wrong, we shall scrap this concept. 

In every phase of science, this practice is followed. Old ideas are 
retained as long as they are useful. They may be altered somewhat 
to fit newly-discovered facts. However, if enough new data are accu- 
mulated to prove them incorrect, the old theories are abandoned. 
It is the belief of scientists that tradition should never be allowed to 
stand in the way of greater enlightenment. 

Albert Einstein, one of the most eminent of living scientists, when 
he was suddenly confronted with new facts which could not fit his 
own theories, expressed the spirit of science thus: "The new facts 
have smashed my old ideas like a hammer blowl" And he went on 
to change some of his most cherished theories. 


Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 289- 
311. Simon and Schuster, New York, 1948. A sketch of Henry 
Moseley, whose lifework was done in 4 short years. Before the 
world knew this genius, he died. 

Moulton, Forest R., and Schifferes, Justus J., Ed. The Auto- 
biography of Science, pp. 502-509. Doubleday Doran & Co., 
New York, 1945. Contains the original papers dealing with the 
origin of the electron theory. 


1 . Electric charges caused by rubbing or friction are called 
static electricity. 

2. Electrons are tiny particles of negative electricity. A 
cathode ray is a stream of electrons. Our knowledge of electrons 
is the result of the work of William Crookes, J. J. Thomson, 
Robert A. Millikan, and others. 

3. Protons are particles oi positive electricity. A proton is 
1837 times heavier than an electron. It was discovered and 
named by Ernest Rutherford. 


4. A neutron is an electrically neutral particle composed of 
one electron and one proton. Its mass is about the same as 
that of a proton, or of the atom of the lightest hydrogen 

5. All atoms are composed of electrons, protons, and neu- 
trons with the exception of the atom of the lightest hydrogen 
isotope which is composed of one electron and one proton. 
Protons and neutrons are located in the nucleus of the atom. 
Electrons revolve about the nucleus in orbits. The relative 
distance between the nucleus and the planetary electrons is 
so great that we may say the atoms of all elements are largely 
empty space. 

6. A fixed number of electrons is required to complete 
each electronic orbit of the atom. The first ring is complete 
when it contains 2 electrons; the second, when it contains 8; 
the third, when it contains 18; the fourth, when it contains 32. 
The outermost ring never contains more than 8 electrons. 

7. The atomic number of any element is the number of 
free protons in the nucleus of its atom. According to the law 
of atomic numbers, the elements can be arranged in a periodic 
table in the order of their increasing atomic numbers. 

8. The atomic weight of any element is the sum of the 
protons and neutrons in the nucleus of one of its atoms. 

9. Isotopes are different forms of the s^me element. They 
possess the same chemical properties, but have slightly different 
physical properties. Isotopes of an element all have the same 
atomic number, but different atomic weights. 

10. The valence of an element is the number of electrons 
that its atom must borrow or lend to complete its outermost 
ring. Electrons in the outermost ring of an atom which may 
be borrowed or lent are valence electrons. 

11. An element is metallic if the outer ring of its atom con- 
tains less than half the number of electrons necessary to com- 
plete the ring. In other words, metals are lenders of electrons. 

12. An element is nonmetallic if the outer ring of its atom 
contains more than half the number of electrons necessary to 
complete the ring. In other words, nonmetals are borrowers 
of electrons. 

13. An element is amphoteric if the outer ring of its atom 
contains exactly half the number of electrons necessary to com- 
plete this outer ring. 

14. An electric current is a flow of electrons. 

15. Chemical activity depends upon the number of elec- 
trons in the outermost ring of the atom of an element. When 
the outermost ring is complete, the element is inert; if the 
outermost ring lends or borrows one electron, the element is 


very active; if the outermost ring lends or borrows three elec- 
trons, the element is not very active. 

16. Chemical union is the shifting or sharing of the electrons 
in the outer rings until a stable condition is reached. 

17. Electrovalent compounds are those formed by the shift- 
ing of single electrons; covalent compounds are those formed 
by the sharing of a pair or pairs of electrons. 

18. Oxidation, a loss of electrons, increases the valence of 
an element. Reduction, a gain of electrons, decreases the 
valence of an element. 

19. New theories are constantly being advanced about the 
nature of matter. Recently it has been suggested that both the 
proton and electron possess wave properties. It is necessary for 
scientists to be ready to give up theories when facts show that 
newer theories are more nearly accurate. 


Group A 

1. (a) How may static electricity be produced? (b) How 
did the discovery of electrons explain this phenomenon? 

2. (a) Who discovered the electron? (b) How did his dis- 
covery explain the behavior of cathode rays? (c) Who deter- 
mined the mass of the electron? (d) What were his findings? 

3. (a) Who discovered the proton? (b) What principle 
of electricity led him to his discovery? (c) Do the protons of 
different elements differ? 

4. (a) What is the neutron? (b) What scientists took part 
in its discovery? (c) Why is the neutron considered approxi- 
mately equal in mass to the proton? 

5. (a) Describe the general arrangement of protons, elec- 
trons, and neutrons in an electrically neutral atom, (b) De- 
scribe the atom in terms of the solar system. 

t . . . 

6. Explain how elements differ according to the electron 

7. (a) Who developed the first periodic table of atomic 
weights? (b) What was learned from arranging the elements 
in such a table? 

8. (a) What relationship between atomic number and pro- 
tons was demonstrated by Moseley? (b) Why is Moseley's 
periodic table more fundamental than Mendeleyeff s? 


9. (a) For the first 17 elements, what is the general re- 
lationship between atomic number and atomic weight? (b) Il- 
lustrate your answer. 

10. (a) Define atomic weight in terms of the electron the- 
ory, (b) Why do most elements have atomic weights which are 
not whole numbers? 

11. (a) What is an isotope? (b) Explain why all isotopes 
of an element have the same chemical properties. 

12. (a) How does the electron theory explain valence? 
(b) How does it explain the behavior of metals? (c) of non- 
metals? (d) of amphoteric elements? 

13. Explain the relationship between the number of elec- 
trons in the outer ring of an atom and the chemical activity 
of that element. 

14. (a) Make a diagram of the structure of the element 
whose atomic number is 20. (b) Describe some of the element's 
chemical properties. 

15. Phosphorus (P) has an atomic weight of 30.98. (a) Dia- 
gram the structure of its atom, (b) Describe its chemical 

16. An atom has a nucleus containing 1? protons and 18 
neutrons, (a) Make a structural diagram of the atom, (b) De- 
scribe its chemical properties. 

17. The atomic weight ot beryllium (Be) is 9.02. (a) What 
are the atomic number, valence, and chemical properties of 
beryllium? (b) Is it a metal or nonmetal? 

18. The atomic weight of curium, one of the newly-discov- 
ered elements, is 242; its atomic number is 96. How many 
neutrons are in the nucleus of one of its atoms? 

19. Make a diagram of (a) the oxygen atom and (b) the 
sulfur atom, (c) Explain, in terms of the electron theory, why 
they resemble one another chemically. 

20. (a) Draw a diagram of the structure of a molecule of 
potassium bromide, KBr. (b) Using your knowledge of the 
electron theory, explain why the two atoms unite. 

21. What is the difference between electrovalence and co- 


22. How does the electron theory explain why two atoms of 
hydrogen unite with one atom of oxygen to form one molecule 
of water? 

23. How does the electron theory explain the fact that cer- 
tain elements are more active chemically than others? 

24. In terms of the electron theory, explain why some com- 
pounds are stable and others unstable. 

25. Explain oxidation and reduction in terms of the electron 
theory. Use the equation for the reducing action of H 2 on CuO 
to make your answer clearer. 

Group B 

26. Study the theories Benjamin Franklin held regarding 
the nature of electricity. Compare his views with the modern 

27. (a) What is a cloud chamber? (b) For what purpose is 
it used? (c) What are alpha particles? (d) fog tracks? 

28. What evidence is there that protons are always found 
inside the nucleus of an atom? 

29. Using the valence-change method, balance the follow- 
ing equations: 

a) HC1 + Mn0 2 - MnCl 2 + H 2 O + C1 2 T 

b) C1 2 + H 2 -> 2HC1 + O T 

c) KC10 3 -> KC1 + 2 T 

d) S + HNO 3 - H 2 SO 4 + NO | 

30. Why does carbon have the 2 valences, -f-2 and -f-4? 

31. Explain why the valence of a free element must be zero. 


1. Harkins predicted the discovery of the neutron, and 
Mendeleyeff predicted the discovery of several elements. Both 
predictions were verified. Such incidents in the history of sci- 
ence are not as rare as in other fields such as economics and 
politics. Give reasons for this. 

2. Construct a model of the sodium atom based on the elec- 
tron theory. You may use wire for the orbits, a small ball for 
the nucleus, and copper rods for the electrons. 



The United States knows that peace- 
ful power from atomic energy is 
no dream of the future. That capa- 
bility, already proved, is here now 
today. I propose an atomic energy 
agency under the United Nations to 
apply atomic energy to the needs of 
agriculture, medicine., and other 
peaceful activities . . . to provide 
abundant electrical energy in the 
power-starved areas of the world. 
Dwight D. Eisenhower, December 8, 

The dream of harnessing the energy of the atom. The atomic 
bomb dropped on Japan on August 6, 1945, was one of those incred- 
ible pinnacles toward which fate drives unsuspecting man. Thou- 
sands of scientists who had studied the atom and discovered the 
electron, proton, and neutron had little idea of the genii they were 
uncorking. Chiefly, they were men and women trying to learn more 
about the nature of matter, each concerned with his own particular 

A few scientists, it is true, saw the possibility of someday learning 
enough about the atom to enable man to release the tremendous 
forces locked within it. If it could be done, the energy locked within 
a single lump of coal would be enough to drive a huge ocean liner 
around the world. As scientists, they hoped for an achievement which 
they believed to be even greater than the discovery of electricity. 
They dreamed of opening a door to an age of limitless power and 
thus lifting the standards of living of all the peoples of the world. 

When the news of triumph finally came, it surprised even the most 
optimistic of scientists. The great marvel, said President Truman, 
". . . is not the size of the enterprise, its secrecy or its cost, but the 




achievement of scientific brains in putting together infinitely com- 
plex pieces of knowledge held by many men in different fields of sci- 
ence into a workable plan." The story of some of these men and the 
bits of knowledge they accumulated has been told in earlier chap- 
ters. More of the story can now be told. 

An amazing scientific discovery startles the world. Late in 1895, 
William Roentgen (rimt'ge?n) was working in a dark room with a 
Crookes tube covered with black paper. He noticed that while an 
electric discharge was passing through the tube, a small screen cov- 
ered with a chemical, barium platinocyanide, lying on a table several 
feet away, gave off a strange glow. 

This was curious and unusual behavior. He thought it must be 
caused by rays, powerful enough to penetrate not only the glass of 
the Crookes tube but several feet of air as well. He tested the pene- 
. trating power of these strange rays by producing them in front of 
several objects of varying hardness, including a hand behind which 
he had placed a sensitized photographic plate. To his astonishment, 
the film when printed showed a hand with the bones much darker 
than the surrounding flesh. He had, so to speak, taken pictures 
through an opaque solid, a truly remarkable feat! 

These rays in some ways acted like light, but differed from light 
in being of much shorter wave lengths which could pass through 
even solid objects. He named these rays x-rays. X-rays are produced 
by the bombardment of matter by a stream of rapidly moving elec- 
trons, or cathode rays. In the cathode tube, as the cathode rays strike 
the anode, they release a small part of the energy as x-rays. 

X-ray machines (right) are used in industry to detect 
flaws inside metal castings like the crankshaft (left). 

ml Klrrtr, 


Another accident, and the curtain rises on the drama of radio- 
activity. Soon after the discovery of x-rays, another accident occurred 
in the laboratory of Henri Becquerel (bek-reT) . He was testing the 
effect of sunlight on various ores, among which he had, fortunately, 
included an ore of uranium. 

Quite by chance he placed a piece of the ore containing the ele- 
ment uranium upon a fresh photographic plate enclosed in a light- 
tight envelope lying on a table in his darkroom. When he examined 
this plate, he found that it had been changed under the very spot 
on which the ore had rested. This was not the sort of accident to 
reach the front page of the newspapers, as the discovery of x-rays had 
done. But its results were tremendously important. 

A new world of radioactivity is discovered. Becquerel could not 
explain what had happened. He repeated the experiment with other 
ores containing uranium. Pitchblende, he found, emitted similar 
rays, which affected a photographic plate even more than the other 
uranium compounds. He suspected some unknown element to be 
the cause, and asked Madame Marie Curie (ku-re') , a Polish girl 
working as a science teacher and research worker in Paris, to under- 
take the isolation of this unknown element. 

Madame Curie and her husband, Pierre, set out to track down the 
cause of this peculiar behavior of pitchblende. They boiled and 
cooked a ton of this ore, sent to them from the pitchblende mines of 
Austria; they filtered and separated out impurity after impurity. 
Years of almost endless work passed, and, though they labored under 
extreme difficulties, Madame Curie wrote years later: "It was in that 

Marie Curie in her laboratory. For her 
isolation of radium and polonium, 
Madame Curie was awarded the 
Nobel prize for chemistry in 191 1. 



miserable shed of a laboratory that we passed the best and happiest 
years of our lives." Finally in 1898 they succeeded in obtaining a few 
crystals of a salt of a new element, which they named radium. 

The new discovery was made public. A strange element had been 
discovered by a woman. Its salts shine in the dark like tiny electric- 
light bulbs and emit heat continuously. This element is a powerful 
poison even acting from a distance. It causes severe burns, and has 
brought premature death to a number of scientists who have han- 
dled it for long periods. It makes the air around it a good conductor 
of electricity. 

Because radium captured the imagination of the entire world, its 
discovery was a great stimulus to further research. This led to the dis- 
covery that a number of other elements resemble radium in their 
ability to break down and emit several kinds of rays. This property, 
called radioactivity,, is possessed by thorium, uranium, polonium, 
radon, and several other elements. 

In 1902, Rutherford and Soddy, another English scientist, ex- 
plained the disintegration or breaking down of radium. Atoms of 
radioactive elements, they said, are not stable. They explode spon- 
taneously, giving off three types of rays: alpha and beta particles and 
gamma rays. The gamma rays are similar to x-rays; the beta rays are 
electrons. Rutherford later showed that the alpha rays are electrified 
particles consisting of nuclei of helium atoms. The disintegration or 
decay of radium is indicated in Table 7. 

What is the half -life of an element? The half -life of an element is 
the time required for the radioactivity of a given amount of an ele- 
ment to decay to half its original value. For example, starting with 
one gram of radium, it takes 1620 years for half of it to change to 
lead. At the end of the next 1620 years, half of the 0.5 gram which 
is left changes to lead, and this process continues at the same rate. 





at. wt. 21 8 

half-life 3 min. 

loses one helium 

nucleus of 

at. wt, 4 and 


at. wt. 214 

half -life f9.5 min.' 

loses one helium 

nucleus of 

at. wt. 4 and 


at. wt. 210 
half-life 140 days 
loses one helium 

nucleus of 

at. wt. 4 and 


at. wt. 206 

a stable element 

which is the end 

product of radium 


~-!^ radiation 

,90$ molecule 

metal cylinder (-) 

tungsten wire ( + ) 9 |ass fube 

to high potential 
to R and detector * 

Fig. 35. Geiger counter. An entering electron pro- 
duces a discharge, causing current to flow in the 
high resistance R until the fall of potential across 
R reduces the potential and discharge stops. The tube 
is then ready for the arrival of a second electron, 
(left) Use of the Geiger counter to test uranium ore 
for the amount of radiation. 

Standard Oil Company (XJ.) 

The Geiger counter, detector of nuclear disintegration products. 

One type of Geiger counter consists of an argon-filled tube contain- 
ing a metal cylinder and a thin wire. Between these two electrodes, 
a very high voltage is maintained and so adjusted that the tube is on 
the verge of discharge. When an electron enters the tube and col- 
lides with a gas molecule, the tube is discharged and a How of cur- 
rent is produced. This current operates a headphone, or produces a 
loud click or flash of light. The greater the radiation, or the closer 
the tube is to the nuclear disintegration, the greater the effect. 

Alpha particles are also detected by the Geiger counter. Neutrons, 
which have no charge, are detected indirectly with the use of a tube 
which contains a boron compound. The boron nucleus absorbs neu- 
trons and produces particles which may be detected. Gamma rays 
are detected by the secondary electrons they produce. 

Alpha particles, which travel at 10,000-20,000 mi. /sec., have the 
least penetrating ability. About five centimeters ot air, a sheet of 
paper, or a thin sheet (0.1 mm.) of aluminum will stop them. Beta 
particles (electrons) , liberated at speeds more than six times that of 
alpha particles, require several meters of air and several millimeters 
of aluminum to absorb them. Gamma rays have still greater pene- 

Fig. 36. Alpha, beta, and gamma rays 
have different penetrating powers and 
are affected differently by a magnetic field. 

radium . , , v 

salt. alpha (a) 

^J_ particles I 

^Hl .....*.. JW"ii 

beta (fiTT"^ 

lead P articles 

box m 9 net 

( foil 






trating power several centimeters of aluminum are required to 
stop them. Fast neutrons have the greatest penetrating power. All 
produce burns. Radiation absorption is measured in roentgens (r.) . 

Uses of radium. Radium, and radon gas sealed in tubes, are used 
in treating skin diseases and cancer. 

Considerable quantities of radium are used in the detection of 
(laws in castings, forgings, and welds. Parts for aircraft and turbine 
casings are among the many kinds of equipment tested in this way. 

The method of testing is simple. From 25 to 1000 milligrams of 
radium sulfate are placed in the center of a circle of articles to be 
tested. X-ray films are placed on the backs of the specimens. The 
penetrating gamma rays from the radium salt produce a shadow- 
graph on the film, quite like the kind obtained with x-rays. Defects 
as small as 0.25 percent of the thickness of the article can be seen 

The price of radium is about $20,000 a gram. Carnotite, an ore of 
both radium and uranium, is found in Utah and Colorado. The rich 
deposits of pitchblende discovered on the shore of remote Great 
Bear Lake in northern Canada in 1931 compete on a favorable basis 
with the huge deposits in the Belgian Congo which constitute our 
most important source of uranium. 

The attack upon the nucleus of the atom. The disintegration of 
radioactive elements indicates that the building blocks of the atom's 
nucleus consist of neutrons and free protons, both of which can be 
emitted. Helium nuclei, which are liberated also by radioactive ele- 
ments, are themselves composed of neutrons and free protons. 

In the effort to learn more about the nucleus, scientists began 
bombarding atoms with various kinds of swiftly moving projectiles. 
One of the most important machines for acceleration of particles is 
the cyclotron developed in 1929 by Ernest O. Lawrence. Some other 

Brookhavcn National Labti 

The magnet in this giant cos- 
matron at Brookhaven has 
an inside diameter of over 
60 feet and weighs 2200 tons. 



accelerators for smashing atoms are die betatron, which * accelerates 
electrons, and the synchroton, cosmotron, and bevatron, all of which 
speed up protons to even greater speeds than the cyclotron. 

During these researches, several discoveries were made the posi- 
tron, artificial radioactivity, and mesons. The positron is an ex- 
tremely short-lived particle, having the same mass as the electron 
but opposite charge. It was discovered and named by Carl D. An- 
derson, of the California Institute of Technology, in 1932. When a 
positron reacts with an electron, both particles disappear and gamma 
radiation is formed. 

In the nucleus of the atom a tremendous energy lies like a coiled 
spring caused, in part, by the many protons which repel each other 
because they are all of the same electric charge. Mesons may be the 
binding energies that hold the protons in the nucleus together. 
There are several types of mesons. The heavier meson (pi-meson) is 
produced when a nucleus is broken up during a collision. The 
lighter meson (raw-meson) is a decay product of the heavy meson. 
Mesons are very short-lived less than a millionth of a second. 

Summary: Units of matter thus far identified. Let us pause tor a 
few minutes to list these units of matter, their mass, charge, and 
other related items of: information. 

'*>''.; ** * ti^M'^ <a 

*-. ./;.|i 

" ^&*"" 


V*** - 

'j^wr""*"* <-* >*~ 

H fc * : * * 





' ^ .*'> t 1 


J (6) 

1/1 837 that of the 
hydrogen atom 



In Crookes tube, as 
cathode rays 

Joseph John 
Thomson, 1 897 

& Proton 


Approximately that 
of hydrogen 



Stripping an electron 
from hydrogen 

Ernest Rutherford, 

2 Neutron 
1 ( nl> 

Approximately that 
of hydrogen 



Bombarding beryllium 
with helium 

James Chadwick, 

1 Positron 
I U t ) 

Same as the 



Radioactive nitrogen 

Carl D. Anderson, 

1 Meson* 

E8. | 

? I 

Heavy, or pi, meson 
is about 285 times 
that of electron. 
Light, or mu, meson 
is about 21 5 times 
that of electron. 


nucleus; as 
cosmic rays 
in upper 

Bombarding atoms 
with primary cosmic 
rays (protons) or with 
helium nuclei of 300 
million or more 
electron-volts in 

Carl D. Anderson 
S. Neddermeyer 
discovered the 
mu-meson in 1 937 
andC. F. Powell 
the pi-meson in 

5 'This particle was predicted by H. Yukawa, a Japanese, in 1934. It was named meson, 
s meaning Intermediate porfic/e It changes into something else in less than a millionth 
$ of ? second, and it is supposed to travel at a speed nearly that of light. 


Ancient and modern alchemy. As you know, modern chemistry 
sprang from alchemy, which was practiced for more than 20 centu- 
ries. The chief goals of the alchemists were to change the base metals, 
such as lead and iron, into gold and to find an elixir that would cure 
all disease. Although among the alchemists there were many honest 
enthusiasts, the annals of their queer practice are filled with accounts 
of liars and charlatans. In many museums of Europe we can still see 
shiny, yellow metals reputed to be gold, made by the deceptions of 
the * 'gold-cooks" of European courts. 

Today the alchemists' dream of changing one element into an- 
other, called transmutation, has come true. Radium changes, of its 
own accord, into helium, lead, and other elements. Besides this natu- 
ral transmutation, however, chemists have succeeded in artificially 
transmuting many nonradioactive elements. The first such transmu- 
tation was achieved in 1919 by Rutherford, who bombarded nitro- 
gen with helium nuclei. The nucleus of the nitrogen atom was 
changed, and one proton was liberated, the remaining nucleus be- 
coming the heavy isotope of oxygen with atomic weight 17. 

How nuclear reactions are written. Transmutation or nuclear re- 
actions are like chemical equations and must balance. The sum of 
the subscripts must be the same on both sides, and so must the sum 
of the superscripts. 

Rutherford's transmutation may be expressed as follows. 

The subscripts (7 + 2-8+l) represent atomic numbers and the 
superscripts (14 + 4 17 + 1) atomic weights. 

Recent experiments indicate that other elements may be built 
up from lighter elements, for example, carbon from beryllium (a rare 
metal lighter than aluminum) . This transmutation, (Fig. 37) , may 
be used to produce high-speed neutrons by mixing powdered beryl- 
lium with a trace of a radium salt which emits helium particles. 

Transmutation, as you see, involves changes in the nuclei of atoms 
rather than in the shifting or sharing of electrons which produces 
only chemical changes. 

Fig. 37. Transmutation of beryllium into carbon. 

Beryllium + Helium ^ Carbon 

(at. wt. (at. wt. 



The search for the key to nuclear energy. Albert Einstein, in 1905, 
advanced the idea that matter and energy were really different forms 
of the same thing, and that matter could be changed into energy 
at least theoretically. He developed a mathematical equation to ex- 
press the conversion of matter into energy: 

E = mc 2 

where E is energy expressed in ergs,* ra is mass expressed in grams, 
and c is the speed of light, expressed in centimeters per second. 

According to this matter-energy conversion equation, one pound 
of matter (for example, one pound of coal or of uranium) is equiva- 
lent to about 11 billion kilowatthours, if completely changed into 
energy. This is about two and one-half times the amount of electric 
energy produced in an entire year by the largest steam-electric gen- 
erating plant in the country. In burning the same amount of coal, 
approximately four kilowatthours of energy are obtained. In terms 
of energy produced, oxidation is, therefore, an extremely inefficient 

These ideas, of course, were all theory. However, a bit of confirma- 
tion came in 1932. In that year, Cockcroft and Walton, working in 
Rutherford's laboratory, bombarded lithium with high-speed pro- 
tons produced by accelerating hydrogen nuclei by Yneans of high 
voltages. They obtained helium (alpha particles) with energies al- 
most 100 times as great as the energy that was used to break the 
lithium atom. This extra energy comes from the conversion of mat- 
ter into energy, in accordance with the equation formulated by Ein- 
stein, thus: 

Lithium + hydrogen > 2 helium + energy 

3 Li 7 + iH 1 2 2 He 4 + 600,000 electron-volts 

Mass 7.0180 + mass 1.0076 -> mass 2(4.0029) f 

8.0256 -> 8.0058 

Approximate loss of mass = 0.02 

* 1 erg == 1/980 gm.-cm. of work = approximately the energy required to lift 
a postage stamp to a height equal to its thickness. The speed of light is 
186,000 mi./sec. or 30,000,000,000 cm./sec. 

f Note that the atomic weight of the isotope of the lithium used differs from 
the atomic weight given in the table on page 162, which is an average of the 
atomic weights of all the isotopes of lithium. The hydrogen here refers to the 
proton, which is slightly lighter than the hydrogen atom, whose weight is given 
in the table on page 162. The weights given in the table on page 162 have been 
rounded off to three decimal places; hence, helium is shown there as 4.003 instead 
of 4.0029 as in this equation. 


However, the method used by these experimenters was not efficient, 
and there was no great excitement over their news. 

The key is found. In the meantime, other scientists were working 
in this same field. In 1934, a young Italian physicist, Enrico Fermi 
(far'mi) , who later left fascist Italy to become professor of physics 
at Columbia University, bombarded uranium with neutrons and 
thought he had created a new element No. 93. Then four years later 
Otto Hahn and F. Strassman repeated Fermi's experiment in Berlin. 
They bombarded uranium with slow neutrons and, instead of pro- 
ducing a new artificial element, they obtained two other natural ele- 
ments and a great deal of energy. Unable to explain what had hap- 
pened, Hahn and Strassman nevertheless published their findings. 

Lise Meitner (mlt'ner) , an eminent woman scientist working with 
Hahn, interpreted the results and passed the information on to Niels 
Bohr, Nobel prize winner in physics, in Copenhagen. Dr. Meitner 
was forced to flee Germany by the Na/is. Dr. Meitner believed that 
when uranium is bombarded by slow neutrons, the atom of uranium 
actually splits by a process called nuclear fission, forming barium and 
krypton. But what is even more important, great quantities of en- 
ergy are released, perhaps as much as 11 million kilowatthours per 
pound of uranium. And, this is only a small part of the energy that 
would be produced if all the uranium were converted into energy. 

The stage is set. Very soon after, a most important conference 
was held in Washington, D.C. Atomic physicists from American col- 
leges and famous scientists from foreign nations were present. Niels 
Bohr was there, and so was Enrico Fermi. At this meeting, Bohr and 
Fermi discussed the ideas of Meitner. Bohr suggested that it was the 
U-235 in the uranium that actually split. Fermi suggested that, in 
the fission of uranium by neutrons, other neutrons might be emitted. 
These emitted neutrons could attack other uranium atoms. If this 
were true, the possibility of a chain, or self-propagating, reaction 
that would unlock the door to nuclear energy was near at last. 

Brown Ilroflu 

Enrico Fermi (1901-1954), winner of the 1938 
Nobel prize for physics, played an important 
role in our government's nuclear research 
program both during and after World War II. 


Before the meeting in Washington was over, experiments to con- 
firm nuclear fission had begun, and confirmation of the emission of 
neutrons was soon obtained. By midsummer of 1940, the important 
facts regarding nuclear fission had been discovered and were known 
by many scientists. And although a chain reaction had not been ob- 
tained, its possibility was clear and several methods of producing 
it had been suggested. Then suddenly, World War II clamped tight 
the door of censorship on all research relating to the release of 
nuclear energy. For five years, the outside world was kept in the dark. 

Nuclear energy unleashed! With the sudden dropping of the first 
atomic bomb on Hiroshima in August, 1945, the veil was partly 
lifted on research on nuclear fission and the production of chain 
reactions. Early in 1940, Franklin D. Roosevelt and Winston 
Churchill had pooled the efforts of British and American scientists 
on a research program, the like of which the world had never seen. 
Its goal was the release of nuclear energy for the production of a 
weapon with which to win the war against the Axis nations more 
quickly. Knowledge that research on such a weapon was being car- 
ried on in Nazi laboratories compelled quick, cooperative action. 
The race was on the prize, the world. The United States gov- 
ernment invested two billion dollars in ". . . the greatest scientific 
gamble in history and won." 

The term atomic energy has long been used to describe the tre- 
mendous power which is released when nuclear fission occurs. How- 
ever, the term nuclear energy is now preferred since it is more truly 
descriptive of the processes involved. 

A chain reaction from U 235. The first controlled chain reaction 
was achieved on December 2, 1942, at the University of Chicago. The 
fissionable material used was pure U-235 obtained from natural 
uranium ores containing a mixture of three isotopes: U-234, U-235, 
and U-238. Even though only about one part in 140 of this mixture 
is pure U-235, this isotope is used because it is most susceptible to 
nuclear, fission by slow neutrons. What happens in the nuclear fission 
of U-235 may be represented as: 

U-235 + neutron - Ba + Kr + 2 or 3 neutrons + energy 

At Oak Ridge, Tennessee, U-235 was laboriously separated from its 
other isotopes by an electromagnetic method. A compound of ura- 
nium, UF 6 , was passed in the form of a gas between the poles of a 
magnet. The lighter isotope, U-235, was deflected more than its 
heavier partners and thus separated. 

liberated neutrons energy 


56 + 
82 n 

Fig. 38. A possible chain reaction. 

A newly created element, plutonium, for the A-bomb. Few details 
about the A-bombs exploded over Hiroshima and Nagasaki have 
been released. It is definitely known, however, that two fissionable 
elements were produced for use in bombs; namely, U-235 and plu- 
tonium. Plutonium, which has fission properties similar to U-235, is 
a newly created element of atomic number 94. It was named for the 
planet Pluto, which lies beyond Uranus in the solar system. Together 
with another new element, neptunium of atomic number 93, plu- 
tonium was first prepared in 1940 with the aid of Lawrence's cyclo- 
tron by E. M. McMillan and P. H. Abelson.* These were momentous 

The formation of neptunium and plutonium may be represented 
by the following three equations: 

1) Uranium 238+ neutron 

uranium 239 

23 mirv 

2) Uranium 239 ^-neptunium 239 + electron 

* Traces of these two elements were later found in some uranium ores. 

In 1945, elements 95, americium (Am) , and 96, curium (Cm) , were obtained 
by bombarding plutonium and uranium with swiftly moving helium nuclei. 

Then came elements 97, berkelium (Bk) , and 98, californium (Cf) , in 1949 
and 1950. These elements were named after Berkeley and California, the city and 
state in which they were first produced. Elements 99 and 100 were created in 1954. 



This change occurs by the breaking down of one neutron in the nu- 
cleus of U-239 into one proton and one electron, which escapes. 

2.3 days 

3) Neptunium 239 *> plutonium 239 + electron 

This change occurs by the breaking down of one neutron in the 
neptunium 239 nucleus into one proton and one electron, which 

Plutonium, in turn, becomes U-235 by natural radioactive disinte- 
gration, or more rapidly by the action of slow neutrons to which it 
is extremely sensitive. The change is indicated by the following nu- 
clear equation: 

Plutonium U-235 -f helium 

What is meant by "critical size"? An A-bomb is set off by suddenly 
bringing together two separate blocks of fissionable material, each 
of which is smaller than the critical size, but which together form a 
mass just above this critical size. For a bomb explosion to occur, the 
number of neutrons captured with fission must be greater than the 
number of neutrons which escape. The number of neutrons which 
escape depends on the surface area, whereas the number captured 
depends upon the volume. As the quantity of fissionable material 
increases, the volume increases faster than the surface area. Critical 
size means the size at which the neutrons captured exceed the num- 
ber which escape and fission occurs. 

Plutonium is produced in nuclear reactors or "piles." The nuclear 
reactor built at Oak Ridge, Tennessee, is essentially a large cube of 
graphite bricks containing a number of horizontal channels into 
which is placed pure uranium in the form of solid cylinders or 
slugs enclosed in aluminum casings. Graphite is used to slow down 
neutrons and is called a moderator. Heavy water is another good mod- 
erator. Slow neutrons are more effective in producing fission than 
are neutrons that travel at normal speeds. The bricks are built up 
in layers, and since the structure was built by piling one layer of 
bricks upon another, it is called anatomic pile. 

A chain reaction is started with neutrons liberated from a bit of 
beryllium mixed with radium in the center of the pile. The concen- 
tration of neutrons is controlled by cadmium or boron-steel rods, 
which absorb neutrons easily. Several nuclear reactions take place as 

concrete shield 
boron steel control rod 

technician removing 
tubes containing 

protective lead shield 

long graphite stringer 

holes for aluminum tubes 

Fig. 39. Simplified drawing of a 
graphite-moderated atomic pile. 

graphite moderator 




Adapted from a drawing of the Atomic Enerf/y Commission 

shown on page 187. When the uranium slugs are ready for processing, 
they are pushed out at the back of the reactor and new ones are fed in 
at the front. The slugs fall into tanks of water where the U-239 grad- 
ually changes into Np-239 and finally into plutonium. The slugs are 
then dissolved in acid, and the plutonium is separated chemically 
from the rest of the elements present. This chemical process of sepa- 
ration is much easier than physical separation of U-235 from U-238. 

Since dangerous radiations and radioactive material are produced 
during these changes, all operations are performed by remote control. 
The whole pile is surrounded with several feet of concrete to shield 
and protect the operating personnel. Several reactors are in opera- 
tion in this country. The one at Oak Ridge is air-cooled, while the 
pile at Hanford, Washington, is water-cooled. Intensive research on 
other coolants is now being carried on. Metals, such as sodium, in a 
liquid state have been found useful for this purpose. 

The atomic pile liberates tremendous amounts of heat. Efforts to 
build industrial nuclear reactors which will utilize this huge source 


Co 60 

Fig. 40. Use of radioactive cobalt for 
the detection of flaws in castings. 

defect in 

Adapted from a drawing of the 
Atomic Energy Commission 


developed film 
shows defect 




of power are already well under way. Atomic power plants "burning" 
nuclear fuel will supply electricity, not only in our own country but 
later on also, it is to be hoped, in those areas of the world which are 
poor in coal, petroleum, and natural water power. Nuclear furnaces 
may be built by private industry with uranium supplied by the 
Atomic Energy Commission. The plutonium manufactured during 
the process will be turned back to the United States Government. 

Radioisotopes, first produced artificially by Madame Curie's 
daughter. In 1934, Irene Joliot-Curie (zho'lyf/) and her husband, 
Frederic Joliot-Curie, by bombarding boron with alpha particles, 
produced a neutron and a radioactive isotope of nitrogen. Here was 
another case of modern alchemy. 

Boron + helium - 
mass 10 + mass 4 - 

> radioactive nitrogen -f neutron 
mass 13 + mass 1 

Since then, scientists have made more than 700 new and different 
radioisotopes in cyclotrons and nuclear reactors. For example, when 
a bit of ordinary iodine (atomic weight 127) is placed in an atomic 
pile where it is bombarded with neutrons, it changes to a radio- 
isotope of iodine of mass 131. The Atomic Energy Commission sup- 
plies hundreds of radioisotopes to research groups all over the world. 

A new "tracer technique" uses radioisotopes. Research in medi- 
cine, biology, agriculture, and many other fields has been helped tre- 
mendously by this new method. Radioactive iodine, for example, is 
being used in thyroid diagnosis and therapy. A person suffering from 
hyperthyroidism is fed with a trace of sodium iodide containing I 131 . 
With the aid of a Geiger counter, the rate at which this iodine com- 
pound collects in the thyroid gland can be accurately determined. 

Radioactive cobalt, Co" , loses radioactivity in five days and is used 

Atomic Ei 

Materials may be made 
radioactive by exposure 
within a nuclear reactor 
such as this water-boiler 
reactor. Such radioisotopes 
have many uses in indus- 
try, agriculture, and medi- 



in cancer therapy as a substitute for radium and x-ray treatments. 
Radioactive carbon, C 1 *, is a wonderful tool in the study of photo- 
synthesis and such problems of human health- as sugar metabolism. 
Radioactive phosphorus, P 32 , is used in agricultural research dealing 
with the accumulation, utilization, and action of phosphate fertiliz- 
ers. Industry is using radioisotopes in the improvement of steel, in 
studying the action of catalysts, in measuring the flow of under- 
ground water, oil, and gas, and in the detection of leaks. 

Nuclear energy in the future? So far, nuclear energy has been 
used mainly as a military weapon, and for research. No one knows 
what the peaceful use of nuclear energy will bring. It seems likely 
that nuclear science will give higher standards of living to all peoples. 

The unlocking of almost unimaginable stores of energy should 
teach man important lessons. Nuclear energy may transform the 
world by improving the health arid raising the standards ot living of 
millions of persons. But this same instrument in the form of a Hy- 
drogen-bomb can destroy civilization as the A-bomb wiped out much 
of Hiroshima and Nagasaki. Therefore, the nations of the world 
must find a way of preventing a war with nuclear weapons from ever 
taking place. 

The H-bomb, based on the fusion of the heavy isotopes of hydro- 
gen into helium and triggered by the 100,000,000C. temperature of 
an A-bomb explosion, could be made of unlimited size. This thermo- 
nuclear reaction may be expressed as follows: 

Deuterium + tritium > helium + neutron -f- energy 
H 2 + H 3 > He + 


The sudden conquest of nuclear energy demonstrated that science 
in a democracy is strong and tremendously creative. By constant 
vigilance, we must strive to keep it so. 

Fig. 41. Simplified drawing of the use of nuclear energy for generating electricity. 

Adapted from a drawing of the Atomic Energy Comnnsifion 

Reactor control console 

Electric power 


A. >*,,. .,. Turbine 

Reactor core 
Uranium rods 







Curie, Eve. Madame Curie. Garden City Publishing Co., 
New York, 1943. In this intimate portrait of her mother, Eve 
Curie tells an entrancing story of the discovery and isolation 
of radium. 

Glasstone, Samuel. Sourcebook on Atomic Energy. D. Van 
Nostrand Company, New York, 1950. This book, written at 
the request of the United States Atomic Energy Commission, 
brings together in readable narrative the important facts about 
the past history, recent status, and possible future of nuclear 

Fermi, Laura. Atoms in the Family. The University of Chi- 
cago Press, Chicago, 1954. A simple, intimate account of the 
events surrounding the conquest of a controlled nuclear chain 
reaction, written by the wife of Enrico Fermi. 

Dean, Gordon. Report on the Atom. Alfred A. Knopf, New 
York, 1953. An account of the role of nuclear energy by the 
former head of the U. S. Atomic Energy Commission. 

Smyth, Henry D. Atomic Energy for Military Purposes. 
Princeton University Press, 1945. This is the so-called Smyth 
Report released by the Army a few days after the first atomic 
bomb was dropped on Japan. It is semitechnical in nature and 
not easy reading. 


1. X-rays consist of vibrations shorter than those of light. 
They are produced by bombarding metals with a stream of 
rapidly moving electrons. These x-rays have great penetrating 

2. Certain elements break down naturally, or disintegrate, 
forming lighter elements including helium nuclei. This prop- 
erty is called radioactivity. 

3. Radium gives off alpha particles, beta particles, and 
gamma rays, which are similar to x-rays. 

4. The simplest units of matter thus far identified are the 
electron, proton, neutron, positron, and meson. 

5. The positron has the same mass as the electron; it has 
a positive charge; and it can be formed during the disintegra- 
tion of radioactive nitrogen. 

6. Mesons have masses between the mass of an electron and 
the mass of a proton; they have either a negative charge, a posi- 
tive charge, or are neutral; they are produced by the bombard- 
ment of atoms with cosmic rays or with rays of 100-million 
electron-volts from the betatron. 


7. Modern alchemy, or artificial transmutation, is an ac- 
complished fact. 

8. Artificially radioactive elements were first obtained by 
the Joliot-Curies by bombarding boron with alpha particles. 
Hundreds of new radioisotopes have since been produced. 

9. The age of nuclear energy was ushered in by the con- 
trolled fission of U-235 and plutonium in 1945. 


Group A 

1 . (a) Who discovered x-rays? (b) How are they produced? 
(c) How do they differ from light? 

2. How did Becquerel's discovery lead to the discovery of 

3. List some of the properties and uses of Ra. 

4. What is meant by the half-life of an element? 

5. What are the alpha, beta, and gamma rays emitted 
during the disintegration of Ra? 

6. Describe the construction and operation of a Geiger 

7. Name five different particles that have been expelled 
from the nucleus of atoms during bombardment. 

8. Name several known facts concerning the mesons. 

9. (a) Has the dream of transmutation come true? (b) Ex- 
plain your answer. 

10. Show by a diagram how Rutherford changed N into H. 

11. (a) Write the Einstein equation for the conversion of 
mass and energy, (b) Illustrate the meaning of this equation 
in terms of the change of lithium into helium. 

12. What is meant by a chain reaction? 

13. How does artificial radioactivity differ from the natural 
radioactivity of Ra? 

14. Describe some of the events since 1938 leading up to the 
final conquest of nuclear energy. 

15. By means of three equations, explain the production of 
plutonium from uranium. 

16. By means of equations, explain how nuclear energy was 
released in 1945. 



17. Describe the construction of a nuclear pile. 

18. By means of an equation, explain the explosion of a 
hydrogen bomb. 

Group B 

19. "The conception of the structure of the atom makes it 
possible for present-day scientists to explain the riddle of 
transmutation." Explain this statement. 

20. Make a diagram of the heaviest known element showing 
the composition of the nucleus and the positions of all its 
electrons. Consult your teacher or a recent edition of some col- 
lege chemistry textbook. 

21. Scientists believe that in releasing about four killowatt- 
hours of energy in burning one pound of coal, a very small part 
of the coal is converted into energy. Why has this fact not been 
proved by experiment? 

22. Describe briefly the future of the peacetime uses of 
nuclear energy. 

23. (a) Would you buy stock in a company organized to 
exploit nuclear energy? (b) Give reasons for your answer. 


1. Make a small model of an electrical power plant utilizing 
nuclear fuel. Consult your bibliography on page 192. 

2. Visit your dentist or doctor and ask him to show you his 
x-ray machine. Make a report on its construction and opera- 
tion, using diagrams. 

3. Consult your teacher of economics on the question: 
"What would the effect of commercial transmutation of iron 
into gold be on the financial structure of the world, if it were 
accomplished tomorrow?" Write a report on this subject. 

4. Compare the effects of the Industrial Revolution and the 
possible effects of use of nuclear energy. Ask your social science 
teacher for help. 

5. Write a two-page report on the construction and opera- 
tion of a cyclotron or other particle accelerator now in use. 




. . . To one man science is a sacred 
goddess to whose service he is happy 
to devote his life; to another she is a 
cow who provides him with butter. 
Liebig, 1803-1882 

Gunfire in the American wilderness helped us to learn more 
about digestion. In. 1822 at a remote fort on Mackinac Island be- 
tween Lake Huron and Lake Michigan, a French-Canadian, Alexis St. 
Martin, was brought in tor medical treatment. An accidentally dis- 
charged musket had sent a bullet through the wall of his stomach. 
Dr. William Beaumont, an army surgeon, patched him up. 

Despite great effort, it was impossible to get the wound to close, 
and on healing, a flap covering an opening into St. Martin's stomach 
was left. Through this opening Beaumont could reach directly into 
St. Martin's stomach. Beaumont got a strange idea. This freak "lid" 
over the hole into St. Martin's stomach would enable him to perform 
experiments to discover the digestive action of the juices of the 
stomach. St. Martin was agreeable and Beaumont tied pieces of food 
to a string, inserted them into St. Martin's stomach and, after several 
hours, removed what was left of the food. 

In this way Beaumont gave science the first accurate facts concern- 
ing the relative digestibility of foods and the composition of gastric 
juice. He found gastric juice to contain a small amount of hydro- 
chloric acid (about 0.3 percent) , which helps to digest certain foods, 
especially proteins. 




Fig. 42. Laboratory prepara- 
tion of hydrogen chloride. Why 
is the end of the delivery tube 
above rather than below the 
level of the water? 

x water 

How hydrogen chloride is prepared in the laboratory. To prepare 
hydrogen chloride in the laboratory, concentrated sulfuric acid is 
added to sodium chloride in a flask, as shown in the illustration 
above. This mixture is heated, and hydrogen chloride gas, which 
is liberated readily, is collected by the displacement of air. In this 
chemical change, a double replacement occurs, as indicated in the 
following equation: 

H 2 SO 4 + NaCl - HC1 + NaHSO 4 (sodium bisulfate) 

At the outset, we must distinguish clearly between HC1 (hydro- 
gen chloride gas) and HC1 (hydrochloric acid) . When hydrogen 
chloride gas is dissolved in water, hydrochloric acid is obtained. 
Both are represented by the same formula, but their physical and 
chemical properties are entirely different. 

How hydrogen chloride is prepared commercially. When a jet of 
hydrogen is burned in chlorine, hydrogen chloride gas is formed (see 
Chapter 10) . 

H 2 + C1 2 -> 2HC1 

This is one commercial method of manufacturing HC1. A second 
commercial method is similar to the laboratory one but at a some- 
what higher temperature producing Na 2 SO 4 instead of NaHSO 4 . 

2NaCl + H 2 SO 4 -> 2HC1 + Na 2 SO 4 

Chemical properties of hydrochloric acid. Hydrochloric acid is 
one of the most common and useful laboratory chemicals, or re- 
agents. Some of its important properties are: 

1) Taste. Dilute hydrochloric acid has a sour taste. 

2) Action on indicators. Hydrochloric acid reacts with a group of 
substances known as indicators, causing a color change. For example, 
hydrochloric acid turns blue litmus pink. It also turns reddish-purple 
phenolphthalein (fe-nol-thal'en) colorless. 


3) Action with metals. When hydrochloric acid is in contact with 
most metals, a reaction takes place. Hydrogen is liberated (recall the 
laboratory method for the preparation of hydrogen) , and chlorides 
are formed. Note that the metal replaces the hydrogen of the HC1. 

2HC1 + Zn -> H 2 t + ZnCl 2 (zinc chloride) 
2HC1 + Fe -> H 2 f + FeCl 2 (ferrous chloride) 

4) Action with bases, or hydroxides of metals. Hydrochloric acid 
reacts with a base, forming a neutral compound that possesses the 
qualities of neither acid nor base (bases are discussed in Chapter 14) . 
Pure water is the only other product of this reaction. For example, 
when hydrochloric acid and sodium hydroxide react, the products 
are common salt and water. 

HC1 + NaOH - NaCl + H 2 O 

Why hydrochloric acid is a typical acid. The chemical proper- 
ties of hydrochloric acid are characteristic of the whole group of com- 
pounds known as acids. We may now define an acid as a water solu- 
tion of a compound with the following characteristic properties: 

1) An acid has a sour taste. The sour taste, or tartness, of fruits is 
caused by certain acids, such as citric acid, which is found in lemons, 
limes, and grapefruit. 

2) An acid turns blue litmus pink, reddish-purple phenolphthalein 
colorless, and acts with other indicators in the same way as hydro- 
chloric acid. 

3) An acid contains hydrogen that can be replaced by most metals, 
forming compounds known as salts. This does not mean that all com- 
pounds containing hydrogen are acids. Sugar, for example, contains 
hydrogen, but it is not an acid, because its hydrogen cannot be re- 
placed by metals. A compound in which the hydrogen of an acid 
has been replaced by a metal is known as the salt of an acid. Thus 
sodium chloride, NaCl, is a salt of hydrochloric acid, and sodium ni- 
trate, NaNO 3 , is a salt of nitric acid, HNO 3 . 

4) An acid neutralizes any base, forming water and a salt whose 
composition depends on both the acid and the base used. 

In general, an acid consists of hydrogen and a nonmetallic ele- 
ment, or hydrogen and a radical. The hydrogen of the acid can be 
replaced by a metal. Strictly speaking, while certain substances show 
acid properties only in water solution, they are commonly called 
acids even when dry. Thus, perfectly dry H 2 SO 4 is called sulfuric 
acid, even though it exhibits acid properties only when it is in a 

Corn Industries Research Foundation, Inc. 

In these converters, corn starch is changed, under heat and pressure, into corn 
sugar, or dextrose. Hydrochloric acid serves as a catalyst in the process. Because 
the process simulates human digestion, dextrose is, in effect, a predigested food 
and is readily assimilated by the body. 

water solution. Bases also exhibit their characteristic basic properties 
only when water is present. Certain characteristics of acids or bases 
in solution are discussed in Chapter 16. 

The general method for the preparation of an acid. Sulfuric acid 
is one of the most important acids, because it is used as a raw material 
in the manufacture of most of the other acids. Sulfuric acid has two 
special characteristics that make it well suited for this purpose: 
(1) its low cost, and (2) its high boiling point (338C.) . The manu- 
facture of sulfuric acid is discussed on page 309. 

The preparation of hydrochloric acid illustrates the general 
method used in preparing acids from sulfuric acid. First, a salt of 
the acid to be prepared is chosen as a source of the nonmetallic ele- 
ment of the acid. Common salt is the least expensive and most abun- 
dant source of chlorine. Many other chlorides can be used, but they 
are more expensive. Sulfuric acid supplies the hydrogen. 

The salt of the acid to be prepared, NaCl, and sulfuric acid are 
heated together gently. Hydrogen chloride gas is produced by the 
reaction and is driven off and dissolved in water in the receiving 
vessel, forming the acid. The higher boiling points of the other com- 
pounds taking part in the reaction prevent their vaporizing and thus 
keep them from passing over into the receiving vessel. 

Many other acids are manufactured by treating their least expen- 
sive and most abundant salts with sulfuric acid. The acid formed is 




usually separated from the reacting substances by methods based on 
the differences in their boiling points. 

Physical properties of hydrochloric acid. Hydrochloric acid is a 
colorless liquid, heavier than water. That is, its specific gravity is 
greater than one. It possesses an irritating odor. Both the boiling- 
point and the specific gravity of hydrochloric acid are determined by 
the weight of hydrogen chloride gas dissolved in the water. Hydro- 
chloric acid containing 20 percent hydrogen chloride gas by weight 
boils at 110C. Impure hydrochloric acid is called muriatic acid and 
is usually yellow in color. In this form it was known for many years 
before Priestley, in 1772, first isolated pure hydrochloric acid. 

Hydrochloric acid cleans metals. Before coating metals, such as 
iron and steel, with plates, films, or coatings of other metals, includ- 
ing chromium, silver, tin, and zinc, the surface of the metal must be 
clean and free of oxides. Removing oxides and otherwise cleaning 
the surface of a metal to be plated or coated is a process known com- 
mercially as pickling. One of the chief industrial uses of hydro- 
chloric acid is in the pickling of metals, especially before coating 
with tin in tinning, or xinc in galvanizing, or with the materials that, 
after firing, result in enamelware. 

Small quantities of hydrochloric acid, usually muriatic acid, are 
used in removing rust stains from vitreous washbasins and lavatories. 
Plumbers often use muriatic acid as a flux before soldering. 

Racks of sheet steel emerge from the pickler. The worker in the foreground is 
dipping the sheets in water to remove the acid. 


Hydrochloric acid is used in making other chemicals. Chlorides 
of many metals, including magnesium chloride, aluminum chloride, 
and zinc chloride, are made by the reaction of hydrochloric acid and 
a carbonate or oxide of the metal. For the most part, such chemicals 
are of very high quality, and are used chiefly by manufacturing chem- 
ists and drug houses, and by druggists. Zinc chloride is used to im- 
pregnate wood to prevent decay, in soldering, and in flame-proofing. 

Silver chloride, AgCl, one of the several light-sensitive silver com- 
pounds used to coat photographic film may be made by the reaction 
of silver nitrate, AgNO 3 , and hydrochloric acid. In the manufacture 
of glucose from starch, hydrochloric acid acts as a catalytic agent. It 
is used also in large quantities in the manufacture of glue and gela- 
tin, in the purification of boneblack, and in the processing of textiles. 

Physical properties of hydrogen chloride gas. Hydrogen chloride 
gas is colorless, heavier than air, and has a sharp, penetrating odor. 
It is extremely soluble in water. If a test tube of hydrogen chloride 
gas is placed mouth downward in water, the water will rise almost 
to the top, as the gas dissolves. Water dissolves about 500 times its 
own volume of this gas under normal laboratory conditions. The 
gas can be liquefied and solidified, just as all other gases. 

Chemical properties of hydrogen chloride gas. Hydrogen chloride 
does not show acid properties unless it is dissolved in water or unless 
some water vapor is present. When dry, it is completely inactive. Its 
attraction for water is so great that it forms a cloud or mist in moist 
air of tiny droplets of hydrogen chloride solution. 

This property may be used as a test for hydrogen chloride gas. If 
you blow across the mouth of a bottle containing concentrated hy- 
drochloric acid, a mist will form. This mist is caused by the moisture 
in your breath combining with the hydrogen chloride vapor that 
rises from the bottle. Hydrogen chloride does not burn. 

Preparation and properties of hydrofluoric acid, HF. The prep- 
aration of hydrofluoric acid follows the general method for making 
an acid. Sulfuric acid reacts with calcium fluoride, CaFo, the most 
common salt of hydrofluoric acid, forming hydrofluoric acid. This is 
the method used by Scheele when he first prepared it in 1771. Cal- 
cium fluoride is the chief constituent of the mineral fluorspar, found 
in several parts of the United States. 

H 2 SO 4 + CaF 2 -> 2HF | + CaSO 4 

Pure anhydrous hydrogen fluoride is a colorless gas which boils at 
room temperature (19.5C.) . It is deadly if inhaled. It dissolves 
in water, forming a colorless acid that vaporizes at low temperatures. 

Calibration marks are carefully 
scratched through the wax coat- 
ing on a graduated glass cylinder 
prior to exposing it to hydrofluoric 

Corning (Hass Works 

Such an acid is called a fuming acid. Since hydrofluoric acid re- 
acts with glass, quartz, and most metals, it is distilled in dishes made 
of lead and must he kept in bottles made of polyethylene or other 
plastics with which it does not react. This acid produces powerful 
burns by poisoning the tissues. 

How glass is etched with hydrofluoric acid. Etching is the biting 
out of particles of glass or metal by means of chemicals. Hydrofluoric 
acid etches glass because it unites with the silicon dioxide, SiO 2 , of 
the glass, forming silicon tetrafluoride, SiF.,, which is a gas. 

SiO 2 + 4HF -> SiF 4 1 + 2H 2 O 

In etching glass articles, such as thermometers, electric-light bulbs, 
and windows, the surface is completely covered with wax and the 
design to be etched is scratched through the wax. The object is then 
brought into contact with the vapor of hydrofluoric acid. When the 
action on the exposed glass has gone as far as necessary, the object is 
removed from contact with the vapor. In frosting the inside of an 
electric-light bulb, a small amount of hydrofluoric acid is poured 
into the bulb, shaken for an instant, and poured out, and the bulb 
is thoroughly rinsed. It is also used in making Freon refrigerants, and 
as a catalyst in the manufacture of high octane gasoline. 

The other halogen acids. Theoretically, both hydrobromic acid. 
HBr, and hydriodic acid, HI, may be prepared by the general method 
used in preparing acids. 

H 2 SO 4 + KBr - HBr | + KHSO 4 
H 2 SO 4 + Nal -> HI j + NaHSO 4 




When first formed they are colorless gases with strong irritating 
odors, but are almost immediately oxidized by the oxygen of the air, 
forming free bromine and free iodine. In water solution, hydro- 
bromic and hydriodic acids are unstable; on exposure to air they 
decompose as a result of oxidation. 

4HBr + O 2 - 2H 2 O + 2Br 2 

Heat of formation and chemical stability. When hydrogen and 
chlorine react, forming hydrogen chloride, a great deal of heat is 
produced. A reaction in which heat is liberated is called an exother- 
mic reaction (ex out, therme heat) . Exothermic reactions con- 
tinue after they are first started by external heat. On the other hand, 
when hydrogen and iodine unite, heat is continuously absorbed and 
heat must be added if the reaction is to continue. A reaction in 
which heat is absorbed is an endothermic reaction. 

The number of calories of heat absorbed or liberated, during the 
formation of a mole (see page 130) of an element or compound, 
is called its heat of formation. A substance which liberates heat dur- 
ing its formation is said to have a positive heat of formation; a sub- 
stance which absorbs heat is said to have a negative heat of forma- 
tion. A compound such as sodium chloride, NaCl, whose positive 
heat of formation is very great (97,800 calories) iswery stable. Hy- 
drobromic acid, HBr, whose positive heat of formation is small (8400 
calories) is somewhat unstable. Hydriodic acid, HI, which has a 
negative heat of formation, is very unstable. 

A knowledge of the heats of formation is very useful to chemists. 
For example, we can tell whether a certain compound will form and 
how easily it can be prepared. A compound formed by replacement 
has a higher positive heat of formation than the compound from 
which it is formed and hence is more stable. 


C1 2 


Br 2 

By similar reasoning we can see why bromine will not replace chlo- 
rine from KC1. 




Clendening, Logan. The Human Body* pp. 74-75. Alfred A. 
Knopf, New York, revised edition. 1945. An amusing story of 
Dr. William Beaumont's experiments with Alexis St. Martin. 

Jaffe, Bernard. Men of Science in America, pp. 157-158. 
Simon and Schuster, New York, 1944. The dramatic story of 
the pioneer investigations of the American doctor, William 
Beaumont, and of John R. Young, which resulted in our 
knowledge of the presence of hydrochloric acid in the gastric 

Lowry, T. M. Historical Introduction to Chemistry, 
pp. 12-16. The Macmillan Co., London, 1915. The discovery 
of the common acids and their chemical and physical proper- 
ties are reviewed. 


1. An acid is a substance whose water solution (1) has a 
sour taste, (2) turns blue litmus pink, (3) contains hydrogen 
that can be replaced by each of many metals with the forma- 
tion of a salt and the liberation of hydrogen, and (4) neu- 
tralizes bases, forming a salt and water. 

2. Chemical reactions in which heat is liberated are called 
exothermic reactions; those in which heat is absorbed are 
called endothermic reactions. 

3. The heat of formation of a compound is the number of 
calories of heat liberated or absorbed in the formation of 
1 gram-molecular weight of the compound. 

4. Compounds with high positive heats of formation are 
stable; compounds with low positive heats of formation are un- 
stable; compounds with negative heats of formation are very 


Group A 

1. Describe the laboratory method of preparing HC1. 

2. (a) What are the two methods of preparing hydrogen 
chloride commercially? Write an equation for each method. 
(b) What type of chemical reaction does each method repre- 


3. Give the properties of HC1 as to (a) taste, (b) action on 
indicators, (c) action on metals, (d) action on oxides, (e) ac- 
tion on hydroxides of metals. 

4. (a) What is a salt? (b) Give three examples. 

5. (a) What are the salts of HC1 called? (b) Name three. 

6. (a) What are salts of HNO 3 called? (b) Name one. 

7. (a) What is an acid? (b) Name five acids. 

8. Make a list of the properties of hydrogen chloride gas. 

9. What is the percentage of hydrogen chloride in a solu- 
tion that has a fixed boiling point of 110C.? 

10. What weight of pure salt is needed to prepare 292 Ib. of 
hydrochloric acid containing 15 percent HC1 by weight? 

1 1 . Compare the solubility of hydrogen chloride in H 2 O 
with the solubility of air, O 2 , H 2 , CO 2 , and N 2 . 

12. What causes the fuming of hydrogen chloride in moist 

13. What test or tests would you use in showing the pres- 
ence of HC1 in a liquid? 

14. (a) What are the uses of HG1? (b) What process in the 
human body depends in part upon HC1? 

15. (a) State the general method of preparing acids. 
(b) Why is H 2 SO 4 so generally used? 

16. Scheele was the first to prepare HF. His method is still 
used today. Describe it. 

17. By means of an equation, explain the etching action 
of HF. 

18. Explain how HF is stored. 

19. Compare the properties of the other halogen acids with 
those of HC1. 

20. (a) What is a calory? (b) What is heat of formation? 

21. (a) What is the difference between an exothermic re- 
action and an endothermic reaction? (b) Give an example of 

22. Explain the relationship between the heat of formation 
and the stability of a substance. 

23. The positive heat of formation of H 2 S is 2730 calories. 
(a) Is H.,S a stable compound? (b) Explain. 


24. Write balanced equations for the following: 

a) Copper oxide + hydrochloric acid * 

b) Hydriodic acid 4- oxygen > 

c) Sulfuric acid + sodium bromide 

Group B 

25. What is the effect of chlorine water on litmus? 

26. When HC1 is boiled, what passes off? Explain. 

27. A bottle of HI turns brown. Explain. 

28. What chemical tests would you use in identifying each 
of the halogen acids? 

29. Why is muriatic acid used in the soldering of metals? 

30. Water may be considered an acid. Explain. 

31. Why is the general method for preparing acids not used 
in the preparation of HI? 

32. In the electrolysis of HC1, the volumes of C1 2 and H 2 
collected are not the same. Explain. 


1. Take home some strips of litmus paper and test various 
foods and other substances found at home for the presence of 
acids. Report your findings to the class. Make a list of common 
foods that react to your test as acids; as bases. What do you 
conclude from your tests? 

2. Prepare a report on the experiments performed by Beau- 
mont with Alexis St. Martin which dealt with the discovery of 
HC1 in gastric juice. 

3. Consider carefully the quotation at the beginning of this 
chapter. Prepare a report for the class on the meaning of this 
quotation. Illustrate with examples from the lives of great 
scientists and inventors. 




. . . Rouelle [the teacher of Lavoi- 
sier], to whom we owe the term base, 
described how natural salts had been 
restricted at first to salts formed by 
the union of acids with alkalies 
which are soluble in water, and im- 
part on the tongue a saline taste. 
T/M. Lowry, 1915 

How sodium hydroxide is prepared in the laboratory. The lab- 
oratory preparation of hydrogen by the action of sodium on water 
was discussed in Chapter 3. Each sodium atom replaces one of the 
atoms of hydrogen in the water molecule. 

2Na + 2HOH -> SNaOH + H 2 1 

After the water evaporates, a white solid, sodium hydroxide, is left. 

Chief properties of sodium hydroxide. Sodium hydroxide is a 

common and useful substance. Some of its important properties are: 

1) Action on indicators. Sodium hydroxide turns pink litmus blue 
and colorless phenolphthalein reddish-purple. 

2) Action with acids. Sodium hydroxide reacts with an acid, form- 
ing a salt, thus causing the original properties of the acid to disap- 
pear along with its own. For example, sodium hydroxide unites with 
hydrochloric acid, forming common salt and water. The presence of 
the salt may be proved by evaporating the water and tasting the 
solid left behind. 

3) Feel. A water solution of sodium hydroxide has a slippery, 
soapy feeling. 



Why sodium hydroxide is a typical base. The chief properties of 
sodium hydroxide are characteristic of the whole group of com- 
pounds known as bases. We may now define a base as a compound 
(1) that contains the hydroxyl group (OH) , (2) whose water solu- 
tion is soapy to the touch, (3) whose water solution turns pink 
litmus blue and colorless phenolphthalein reddish-purple, and (4) 
whose water solution reacts with acids, forming water and a salt of 
the acid used. 

In general, then, a base consists of a metallic element or radical and 
one or more hydroxyl groups (OH) . 

Strong bases, such as sodium hydroxide and potassium hydroxide, 
are often known as alkalies. Three other important bases are cal- 
cium hydroxide (slaked lime) , ammonium hydroxide (ammonia 
water) , and magnesium hydroxide (milk of magnesia) . 

How bases react with acids. When a base reacts with an acid, the 
hydrogen of the acid combines with the hydroxyl radical of the 
base, forming water, and the metal or metallic radical of the base 
unites with the nonmetal or acid radical of the acid, forming a salt. 
Such a chemical change is called neutralization. All neutralization 
reactions are double replacements. 

HC1 + NaOH - HOH + NaCl (a salt) 
HNOa + KOH -* HOH + KNO 3 (a salt) 
H 2 SO 4 + Ca(OH) 2 - 2HOH + CaSO 4 (a salt) 
HBr + NH 4 OH -> HOH + NH 4 Br (a salt) 

Neutralization is the reaction of an acid with a base, forming wa- 
ter and a salt. A salt, then, is a compound made up of a metal or a 
metallic radical and a nonmetal or an acid radical. 

Titration and the use of molar and normal solutions. It is some- 
times necessary to know how much of an acid, a base, or a salt is 
present in a solution. For example, we may want to know how much 
acetic acid is present in a given sample of vinegar. One way is to 
neutralize a given volume of this liquid with a solution of a base 
whose composition is known. This process of determining the 
strength of an acid or a base with the help of neutralization reaction 
is called titration. It is carried out in long tubes (burettes) which 
are measured off in milliliters so that volumes can be read off di- 
rectly. A definite volume of the acid solution to be tested is neutral- 
ized with a standard solution of a base, that is, one whose composi- 
tion is known. The point at which neutralization is reached (end 
point) is determined by the use of an indicator such as phenolphtha- 
lein which suddenly changes color. 



Fig. 43. Titration. The degree of acidity of 
any acid solution is determined by measur- 
ing the amount of normal basic solution 
required to neutralize it. The process is 
reversed to find the strength of any basic 

-* burette 


t T *%&(* 

* -" 

There are two kinds of standard solutions used, namely, molar 
solutions and normal solutions. One liter of a molar solution con- 
tains one mole or gram-molecular weight of dissolved substance. For 
example, a molar solution of Nad contains 23 + 35.5, or 58.5, grams 
of NaCl in one liter of this solution. A molar solution of Ca (OH) 2 
contains 74 g. 

One liter of a normal acid solution contains one gram of replace- 
able hydrogen. One liter of a normal basic solution contains 17 grams 
of the OH group, and one liter of a normal salt solution contains the 
equivalent of one gram of replaceable hydrogen. For example, one 
liter of a normal solution of H 2 SO 4 contains 98 -4- 2, or 49, grams 
H 2 SO 4 since this acid contains two replaceable hydrogen atoms. One 
liter of a normal Ca (OH) 2 solution contains 74 -:- 2 grams Ca (OH) 2 . 
One liter of a normal A1C1 8 solution contains 133.5 -4- 3 grams A1C1 S 
since this salt contains three hydrogen equivalents. Decimal frac- 
tions are used in referring to both molar (M) and normal (N) solu- 
tions. Thus a 0.1M solution of HC1 contains 36.5 -~ 10, or 3.65, 
grams HC1 per liter of solution. 


1. How much NaCl per liter of solution does a 0.3M solu- 
tion of NaCl contain? 

2. How much acid, base, or salt per liter of solution do the 
following solutions contain: 0.1N AgNO 3 , 0.3N MgBr 2 , 
1.5N H 3 PO 4 , 0.125N CuCl 2 ? 

3. How would you prepare the following solutions: 

0.8N HNO 3 , 3.0N KOH, 0.25N Mg(OH), 

1.3N A1(OH) 3 ? 


Heat of neutralization. All neutralization reactions are exothermic. 
When any strong acid and strong base react, a definite quantity of 
heat is liberated, 13,700 calories per gram-molecular weight, or mole, 
of water formed. This heat is called heat of neutralization. When 
36.5 grams of hydrochloric acid react with 40 grams of sodium hy- 
droxide, 18 grams of water and 58.5 grams of sodium chloride are 
formed with the liberation of 13,700 calories. In the case of the ac- 
tion of sulfuric acid on potassium hydroxide, 2 X 13,700 calories are 
produced, because two moles of water are formed, 

H 2 SO 4 + 2KOH - K 2 SO 4 + 2H 2 O 

Important uses of neutralization. After petroleum has been puri- 
fied with the aid of sulfuric acid, the excess acid is removed by neu- 
tralizing it with a base, usually sodium hydroxide. Much of the 
soil in the United States is sour (acid) . Four-fifths of the cultivated 
land in the central western states is sour and therefore not fully 
productive. These acid soils may be neutralized by the addition of 
lime, CaO, which combines with water, forming a base. Calcium car- 
bonate, CaCO a , is used also. Soil with excess lime is neutralized with 
ammonium sulfate, (NH 4 ) 2 SO 4 . 

The destructive effect of acid stains or burns may be minimized 
by the prompt application of either a weak base or sodium bicar- 
bonate, which neutralizes the effect of the acid. Similarly, an alkali 
burn is treated by application of a mild acid, such as boric acid or 
vinegar. An excessive acid mouth condition is sometimes treated by 
using small amounts of milk of magnesia, Mg (OH) 2 , a mild base, or 

Normal blood is slightly alkaline. Alkaline reserve refers to the 
amount of base present as bicarbonate in the blood. Even slight 
changes in the normal alkalinity of the blood result in serious body 
disturbances. Such disturbances may, in some cases, be corrected by 
using neutralizing chemicals administered by a physician. 

Existence of an alkaline reserve (a scientific fact) is, unfortunately, 
used by some advertisers as a reason for selling to the public huge 
quantities of "alkalizers" to dose ailments that are best treated by 
other methods administered by a competent physician. This is an 
excellent example of the abuse of scientific knowledge, against which 
intelligent persons must always be on guard. 

Physical properties of sodium hydroxide. The white solid, sodium 
hydroxide, is very soluble in water. On exposure to moist air, it ab- 
sorbs large quantities of water and changes into a pool of sodium 
hydroxide solution. This property of deliquescence makes sodium 

hydroxide useful as a drying agent. Usually it is sold in pellets that 
must be kept in well-sealed bottles which exclude moisture. 

Chemical properties of sodium hydroxide. In addition to typical 
properties of bases already mentioned, sodium hydroxide has other 
chemical properties. When exposed to air, it unites with carbon di- 
oxide, forming sodium carbonate. 

2NaOH + CO 2 -> Na 2 CO 3 + H 2 O 

The common name for sodium hydroxide is caustic soda, given to 
it because of its caustic, or burning, action upon the skin. It dis- 
solves wool but has little effect upon cotton, rayons, or nylons. Be- 
cause of this fact, it is used in determining the amount of wool in a 
cotton-wool mixture. 

Potassium hydroxide (caustic potash} closely resembles sodium hy- 
droxide in both chemical and physical properties. 

Sodium hydroxide helps clothe us. Sodium hydroxide is one of 
the most useful compounds known. It serves man in many ways, 
chief among which is in the manufacture of rayon and cellulose 
films similar to cellophane. Approximately one-fifth of the 3 million 
tons of sodium hydroxide produced in a recent normal year in this 
country was consumed in this way. The chemistry of the production 
of rayon and cellulose films is discussed in detail in Chapter 37. 

Many fabrics are made of cotton that has been treated with so- 
dium hydroxide. When cotton fibers are placed in a solution of 
sodium hydroxide, they lose part of their natural twist and acquire 
a gloss that is considered very desirable by many persons. Cotton so 
treated is known as mercerized cotton. 

Sodium hydroxide helps keep us clean. Until very recently sodium 
hydroxide and potassium hydroxide were used in making almost all 
the cleansing agents for both industrial and home use. When sodium 


hydroxide or potassium hydroxide reacts with a fat, soap and glycerin 
are formed. Each year the soap industry uses thousands of tons of 
sodium hydroxide and considerable amounts of potassium hydroxide 

Another important use of sodium hydroxide is in the form of he 
which contains about 94 percent sodium hydroxide. It is a useful 
household cleansing agent because it dissolves grease. Lye produces 
pamlul burns if it comes in contact with the skin. It should be used 
cautiously and stored out of the reach of small children 

When kitchen or bathroom plumbing becomes clogged, a strong 
solution of lye is sometimes poured down the pipes. In this way 
greasy, fatty accumulations are saponified and become soluble in wa- 
ter. "Drano" and other common plumbing cleaners sold under vari- 
ous trademarks contain percentages of impure sodium hydroxide 

Sodium hydroxide helps peel fruits and vegetables. Many fruits 
and vegetables are peeled before canning or dehydration by the use 
of sodium hydroxide, or lye. Most of the large lye peelers are of the 
moving conveyor type, and the fruits or vegetables are immersed in a 
20- to 25-percent lye solution for from two to five minutes. Durino 
this time, the lye solution reacts with the pectins, or binding mate* 
rials, between the individual cells. The skins become loose and may 
be removed by washing which also removes all traces of lye. 

This process is an outgrowth of the making of old-fashioned lye 
hominy, a staple in the diet of American pioneers. In the making of 
lye hominy, grains of corn are soaked in a lye solution until the hard, 
tough skin of the corn grain becomes loosened. Washing with fresh 
water removes the skin and also the lye. 

Other uses of sodium hydroxide. Large quantities of sodium hy- 
droxide are used in reclaiming aluminum and salvaged rubber, in 
the processing of many vegetable oils, and in the production of gas- 
oline. In the refining of petroleum, large quantities of this "basic 




heavy chemical" are used to neutralize the sulfuric acid with which 
petroleum is purified. Of only slightly less importance is the use of 
sodium hydroxide in digesting and purifying the cellulose of wood 
pulp that is used in manufacturing paper. Potassium hydroxide is 
the electrolyte in certain types of storage batteries such as the Edison, 
and Ni-Cd batteries. Lithium hydroxide is used in submarines to 
absorb CO 2 . 

How sodium hydroxide is prepared for industrial use. Chlorine, 
hydrogen, and sodium hydroxide are all formed at the same time dur- 
ing the electrolysis of brine (see Chapter 10) . Both chlorine and so- 
dium hydroxide are prepared industrially by this method. If chlo- 
rine is the chief product desired, then sodium hydroxide is the 
byproduct. That is, it is a substance formed incidentally during the 
preparation of another substance. If sodium hydroxide is the com- 
pound being manufactured, then chlorine is the byproduct. 

The apparatus that is commonly used in the industrial prepara- 
tion of sodium hydroxide by the electrolytic process is the Hooker 
cell. The Nelson diaphragm cell, Vorce cell, and mercury cell are 
also used. The graphite anodes (see Fig. 44) are covered by a so- 
dium chloride solution and are separated from the cathode by an 
asbestos diaphragm, which prevents the chlorine from mixing with 
the sodium hydroxide. Chlorine gas escapes through an outlet in the 
dome at the top of the cell, and hydrogen gas passes through an out- 
let from the steel screen cathode. Because the sodium hydroxide 
solution is heavier than the salt solution, it concentrates with it at 
the bottom of the cell. It is drawn off, and evaporated to dryness 
during which almost all of the NaCl crystallizes out. This process 

Rased on a diagram by 

Hooker Electrochemical Company 




graphite anode 


Fig. 44. Hooker cell used in the pro- 
duction of sodium hydroxide and 
chlorine. Hydrogen is a byproduct. 


is continuous, and more brine is added as the strength of the sodium 
chloride solution diminishes. 

2NaCl + 2H 2 O - 2NaOH + H 2 1 + C1 2 1 

An older method, still very widely used, depends upon the conver- 
sion of the cheaper base, calcium hydroxide, into sodium hydroxide 
by means of a solution of sodium carbonate. Until 1940, more so- 
dium hydroxide was produced by this process than by electrolysis. 

Ca(OH) 2 + Na 2 CO 3 - CaCO 3 [ + 2NaOH 

Since calcium carbonate is insoluble, it is separated from the soluble 
sodium hydroxide by filtration. 

Methods of preparing a salt. Most inorganic compounds, such as 
sodium chloride, sodium nitrate, copper sulfate, and so forth, are 
salts. We have already had occasion to refer to six of the seven meth- 
ods of preparing salts. A list of these seven methods follows: 

1) Neutralization: 

HC1 + NaOH - H 2 O + NaCl 

2) Action of an acid on a metal: 

3) Union of a metal and a nonmetal: 

Fe-f S-+FeS 

4) Action of an acid on the oxide of a metal: 

2HC1 + CuO -> H 2 O + CuCl 2 

5) Action of an acid on a salt of a more volatile acid: 

H 2 SO 4 + 2NaCl -> 2HC1 + Na 2 SO 4 

6) Action of one salt on another salt: 

AgNO 3 + NaCl -> AgCl + NaNO 3 

7) Action of the oxide of a metal (basic oxide) on the oxide of 
a nonmetal (acidic oxide) : 

CaO + SO 8 - CaSO 4 
The horizontal lines under certain formulas indicate salts. 



Fabre, Jean H. The Wonder Book of Chemistry, pp. 154- 
170. Albert & Charles Boni, New York, 1922. Salts and neu- 
tralization are discussed in a captivating manner. 

Jaffe, Bernard. Chemical Calculations, pp. 96-102. World 
Book Co., Yonkers, New York, 1947. Normal and molar solu- 
tions and problems involving them are included. 


1. A base is a substance that contains a metal or metallic 
radical and one or more hydroxyl groups. Its water solution 
is soapy to the touch, turns pink litmus blue, and reacts with 
acids, forming water and a salt. 

2. Neutralization is the action of an acid with a base, form- 
ing water and a salt. The hydrogen of the acid unites with the 
hydroxyl radical of the base, forming water. 

3. When a strong acid and a strong base react, forming a 
mole of water, 13,700 calories of heat are liberated. This 
amount of heat is called heat of neutralization. 

4. A salt is a compound made up of a metal or a metallic 
radical and a nonmetal or an acid radical. 

5. The seven methods of. preparing salts ar^: (1) neutraliza- 
tion, (2) action of an acid on a metal, (3) union of a metal 
and a nonmetal, (4) action of an acid on a metallic oxide, 
(5) action of an acid on a salt of a more volatile acid, (6) re- 
action between two salts, and (7) reaction of a basic oxide on 
an acidic oxide. 


Group A 

1. (a) Write an equation illustrating the method of pre- 
paring NaOH in the laboratory, (b) How can the NaOH 
formed be changed into a solid? 

2. (a) What other base can be made in the same way? 

(b) Write the equation for the preparation of this other base. 

(c) What is its common name? 

3. (a) What is a base? (b) List the four properties by 
which a base can be recognized. 

4. Give the names and formulas of three bases other than 
NaOH that are often found at home. 

5. (a) What is a salt? (b) What is neutralization? (c) How 
is mercerized cotton prepared? 



6. A soil is found to be acid. To obtain the best results 
with certain crops a neutral soil is required. How would a 
farmer correct the acid condition of his soil? 

0.5M AlBr 3 , and 

7. Write four equations illustrating neutralization. 

8. (a) What is a molar solution? (b) What weight of acid, 
base, or salt do the following solutions contain: 

(1) 3M NaOH, (2) 0.1M H 2 SO 4 (3) 
(4) 1.8M HN0 3 ? 

9. (a) What is a normal solution? (b) What weight of 
acid, base, or salt do the following solutions contain: 

(1) 0.1N Mg(OH) 2 , (2) 0.5N H,PO 4 , (3) 2.5N AgNO 3 , 
and (4) 1.5N A1C1 3 ? 

10. What is meant by titration? 


11. Describe the electrolytic process for making NaOH, 
illustrating your description with a diagram. 

12. In a process of the kind mentioned in exercise 11, what 
determines which is the product and which the byproduct? 

13. Write the equation for a method of preparing sodium 
hydroxide other than the electrolytic method. 

14. (a) What is the annual consumption of NaOH in the 
United States? (b) What are the three chief uses of NaOH? 

15. (a) What is lye? (b) Why is it useful in the home? 
(c) What cautions should be exercised in its use? 

16. A sample of fabric was placed in a test tube containing 
a solution of NaOH, and boiled for a few minutes. Half of the 
fabric disappeared. What can you say about its composition? 

1 7. Copy and complete the following table, inserting the cor- 
rect formulas for the salts formed. Do not write in this book. 



H 2 SO 4 

H 2 C0 8 



Ca(OH) 2 

NH 4 OH 

(NH 4 ) 2 C0 8 

18. Classify as acids, bases, or salts: A1(OH) 8 , K 2 CO 8 
CuSO,, Pb (OH) HC 2 H 8 2 , CaCl 2 , BaSO 4 , HCN, H 2 S. 


19. Which will produce a greater weight of alkali, 45 g. of 
Na or 50 g. of K? 

20. A piece of Na is thrown in H 2 O. The NaOH formed 
exactly neutralizes 24.5 g. of H 2 SO 4 . What weight of Na was 

Group B 

21. (a) Write the equations for the neutralization of sodium 
hydroxide with (1) hydrochloric acid, and (2) sulfuric acid, 
(b) The heat of neutralization for the first reaction is 13,700 
calories; for the second reaction, 2 X 13,700 calories. Inspect 
the two equations. Explain why the heat of neutralization is 
twice as large in the second reaction. 

22. What volume of a 0.1N KOH solution is needed to 
neutralize 50 ml. of a 0.5N solution of H 2 SO 4 ? 

23. What is the reason for adding lime water to milk that is 
to be fed to an infant? 

24. NaHCO 3 is frequently used to neutralize an acid, (a) Is 
it a base? (b) Explain your answer. 

25. A glass bottle with a glass stopper contained solid 
NaOH. When an attempt was made to open the bottle, it was 
found that the glass stopper was firmly cemented to the neck 
of the bottle. Explain. 

26. In making cream of tomato soup, a pinch of NaHCO 3 
is added to the tomato puree before the milk is added. Explain 
the reason for this practice. 


1. Take home some pieces of litmus paper and test the fol- 
lowing substances for acidic, basic, or neutral properties: vine- 
gar, milk of magnesia, ammonia water, table salt solution, 
soap solution, sugar solution, and the liquid in an automobile 
battery. What do you conclude about the importance of acids 
and bases in everyday living? 

2. With the help of your chemistry teacher, perform the fol- 
lowing experiment: Pour an excess of NH 4 OH on a solution 
of FeCl 3 . You will get a brown, sticky precipitate of Fe (OH) 3 . 
Boil thoroughly, then filter. Wash the insoluble Fe (OH) 3 
several times with distilled water. What properties of soluble 
bases are not possessed by this insoluble base? 

3. Using litmus paper as indicator, neutralize (with con- 
stant stirring) a known volume of vinegar with some house- 
hold ammonia. Report to your class on the relative strength 
of these solutions. 




. . . The sea is the chemist that dis- 
solves the mountains and tlie rocks, 
pulverizes old continents and builds 
new, forever redistributing the solid 
matter of the globe. Ralph Waldo 
Emerson, 1803-1882 

Is chemically pure water found in nature? The purest form of 
water in nature is rain water. Although it is perfectly safe for drink- 
ing purposes, rain water is not pure water, for it is mixed with the 
gases of the atmosphere and with small amounts of dust and other 
impurities, which it has washed down from the air. 

Eventually all water finds its way to the ocean, the great reservoir. 
Rain may fall upon busy streets and flow through sewers to' the sea. 
It may fall on the ground and furnish the water so necessary to grow- 
ing plants. Water evaporates from the leaf-surfaces of plants and 
from the surfaces of streams, lakes, and oceans only to fall again as 
rain. Without this water-cycle, life as we know it could not exist. 

Water that flows over the ground (surface water} collects fine 
particles of solid material. The size and the weight of this material 
depend on the speed with which the water flows, for fast-moving wa- 
ter can carry much heavier material than a slow-moving stream. 

Water that soaks into the ground (ground water] carries almost no 
load of this type, since the soil acts as a filter and holds the solid par- 
ticles back. But ground water is still not pure water, no matter how 
clear and sparkling it may be, for it contains minerals that have been 
dissolved out of the soil. 



What is a solution? When a teaspoonful of sugar is dissolved in a 
glass of water, the sugar completely disappears from view, and we 
say that we have made a solution of sugar and water. The substance 
that is dissolved is called the solute; the substance in which the 
solute dissolves is called the solvent. Thus, when we make a solution 
of sugar and water, the sugar is the solute and the water is the solvent. 

If we taste a sample of this solution of sugar and water, we find that 
it tastes the same, no matter whether we have taken our sample from 
the top or the bottom of the solution. This indicates that a solution 
has the same composition throughout; that is, it is homogeneous. 
If we take care. not to let any of the water evaporate, the sugar will 
not settle to the bottom; the solution will remain the same. Of course, 
different amounts of sugar will dissolve in a glass of water the 
solution may be very sweet or it may be only slightly sweet. Thus 
a solution differs from a compound, for we can vary the composi- 
tion of a solution. The composition of a compound does not vary. 

A uniform mixture of solute and solvent that does not conform to 
the law of definite proportions is called a solution. The kinetic 
theory of matter helps to explain the mechanism of solution. Accord- 
ing to this theory, the solute breaks down into molecules which dis- 
tribute themselves between the molecules of the solvent. For this rea- 
son, a solution is sometimes called a molecular dispersion of a solute 
in a solvent. If the solute is colorless (for example, sugar) , it can no 
longer be identified by sight when it is in solution. 

Difference between dilute and concentrated solutions. A pinch of 
salt in a gallon of water makes a very dilute solution. Half a pound 
of salt added to the same amount of water makes a concentrated solu- 
tion. When only a small amount of solute is dissolved in a solvent, we 
have a dilute solution; when the amount of solute dissolved is con- 
siderable we have a concentrated solution. 

What determines the amount of solute that will dissolve? There 
is a limit, of course, to the amount of a solute that will dissolve in 
a given volume of a solvent. Several factors determine the amount of 
solute that will enter into solution. The most important factors are 
the nature of the solvent and the nature of the solute. Large amounts 
of certain substances, such as salt, dissolve in water, but the amount 
of gold that dissolves in water is extremely minute. Iodine dissolves 
only slightly in water, but it is very soluble in alcohol, forming an 
alcohol solution known as a tincture. Just why a substance dissolves 
in one solvent and not in another is not thoroughly understood. 

Temperature has a great deal to do with the amount of a solute 
that will enter into solution. More sugar will dissolve in hot tea than 



in iced tea. In general, most solids dissolve in larger amount in warm 
liquids than in cold liquids. This is not true of gases, for the higher 
the temperature of a liquid, the smaller the amount of gas that will 
dissolve in it. You are familiar with this fact if you have ever noticed 
that gas from a warm bottle of soda pop escapes more rapidly than gas 
from an ice-cold bottle of the same beverage. 



Solubility varies with ^ 
I temperature. A nearly ^ * 
] straight line indicates H 
I that an increase in 
I temperature produces 
1 a regular increase in 

10 20 

30 40 50 60 70 
Temperature, C. 

Another important factor that determines the solubility of a gas is 
pressure. The higher the pressure, the greater the amount of gas that 
will dissolve. For this reason, when we want to make a concentrated 
solution of carbon dioxide in water (the carbonated water you see at 
every soda fountain) , we add the gas to the water under high pres- 
sure and at low temperature. The weight of gas that dissolves is pro- 
portional to the pressure on the gas (Henry's law) . In other words, 
if we double the pressure, twice as much gas will dissolve. 

Two factors that help to determine the amount of solute that will 
dissolve in a definite period of time (speed of solution) are the degree 
of subdivision of the solute and the extent to which particles of solute 
and solvent are closely intermingled by stirring. As you know, the 
finer the particles and the more vigorous the stirring, the quicker the 
solid will dissolve. However, while these two factors affect the speed 
at which a solute dissolves, they do not affect the maximum amount 
of solute that will dissolve. 

What is a saturated solution? If we add a salt, such as alum, to a 
given volume of water at a definite temperature and under fixed 

small crystal 
of hypo 

solution of hypo 

(sodium thiosulfate) 

of hypo 


crystals ' 
starting to form 

hypo crystals 
Fig. 45. Formation of hypo crystals from a supersaturated solution 


conditions of pressure, the water will continue to dissolve the salt un- 
til it has dissolved a certain amount. After this amount is dissolved, 
no more alum will dissolve, and any additional alum that is added 
will settle to the bottom of the water and remain there. A solution in 
which no more of the solute will dissolve, at that particular tempera- 
ture and pressure, is a saturated solution. As long as the solvent will 
dissolve more solute, the solution is said to be unsaturated. 

The addition of another crystal of the solute to a solution will 
indicate saturation or unsaturation. If the crystal added does not 
dissolve, the solution is saturated. From the point of view of the 
kinetic theory of matter, some of the crystals added to a saturated 
solution do dissolve, but just as many molecules of the solute come 
out of solution. Thus the crystals added do not appear to dissolve. 

How supersaturation is used in purifying solids. It may sound 
contradictory to say that it is possible to prepare a supersaturated 
solution, that is, a solution which contains more of the solute than a 
saturated solution. But it is possible. For example, to prepare a super- 
saturated solution of hypo (Na 2 S,O 8 - 5H,O) , we first make a satu- 
rated solution of this salt in boiling water, and then slowly cool the 
solution. The excess hypo does not come out of solution, as we would 
expect, but remains in solution. Because this solution contains more 
solute than it normally holds at the lower temperature when satu- 
rated, it is called a supersaturated solution. 

Supersaturation is an unstable condition, and if the solution is 
disturbed by adding a tiny crystal of the solute, all of the excess salt 
separates out and the solution becomes saturated. Since only pure 



crystals will separate out, this method is often used commercially in 
preparing chemically pure (C.P.) salt crystals. Because impurities 
are present to a small degree as a rule, they separate out later. 

The effect of a solute on freezing and boiling points. As you know, 
the addition of salt to water lowers the freezing point of the solu- 
tion. In general, a solute raises the boiling point and lowers the 
freezing point of a solvent a definite number of degrees. This prin- 
ciple is used in preparing antifreeze mixtures, which freeze at tem- 
peratures much lower than 0C., the freezing point of pure water. 
A water solution that is 50 percent grain alcohol does not freeze un- 
til it reaches 31C. Hence, solutions containing alcohol or other 
solutes are used in automobile radiators in cold weather. 

All substances do not have the same effect on freezing and boiling 
points. Certain acids, bases, and salts raise the boiling point and 
lower the freezing point two or three times as much as other sub- 
stances including sugar, alcohol, and glycerin. This singular behavior 
has an important relation to dissociation (see Chapter 16) . 

The difference between a solution and a suspension. If a substance 
does not actually dissolve but becomes mixed somewhat uniformly 
with a solvent and then separates on standing, the mixture is not 
a true solution but a suspension. Fine clay or other fine materials 
intimately mixed with water is a suspension. 

The chief differences in behavior between a true solution and a 
suspension depend upon differences in the size of the particles of the 
solute in solution compared with the size of the particles in suspen- 
sion. Particles of solutes in solution are in the molecular state of 
subdivision. But even the finest particles in suspension are much 
larger than molecules; hence, they settle out. 

Particles in suspension may be separated by filtration; those in a 
solute cannot be separated by filtration. The solute of a true solution 
has a greater effect upon the boiling and freezing points of the 
solvent than the particles in suspension have on a suspension. 

In the borderland between true solutions and suspensions is an- 
other condition of matter called the colloidal state. This condition 
of matter is so important that all of Chapter 38 is devoted to it. 

How we purify water. Even before the germ theory of disease 
was proved, several methods were in use for the purification of 
water. In ancient times, water was made more fit to drink by filter- 
ing, boiling, or by allowing suspended impurities to settle out. Laws 
were passed to guard against the contamination of river water used 
for drinking purposes by prohibiting the washing of clothes and the 
disposal of refuse and sewage in it. 


In a chemical sense most drinking water is not pure, tor it con- 
tains dissolved minerals and gases. It is pure in a biological sense, 
which means that it is comparatively tree from bacteria and other 
organic matter harmful to health. 

Water used for drinking purposes may be put through one or 
more of the following processes of purification, depending upon the 
nature of the impurities it contains: 

(1) Aeration consists of spraying the water into the air or letting 
it tlow down a series of steps in thin sheets so that sunlight and oxy- 
gen may kill most bacteria present. (2) Filtration consists of strain- 
ing the water through a suitable sieve (filter) , thus separating sub- 
stances either in suspension or afloat. Sand filtration dates as far back 
as 1829, when London purified its drinking water by passing the 
water through beds of line sand. (3) Chlorination, discussed in 
Chapter 10. (4) Ozonalion, discussed in Chapter 2. (5) Coagulation 
consists of adding alum or other chemicals that cause the organic 
matter containing bacteria and other suspended material to settle 
out. Coagulation is discussed more fully on page 396. (6) Chlor- 
amination, another process of water purification using both chlo- 
rine and ammonia is discussed on page 253. (7) Demineralization 
by means of synthetic resins (see page 191) . 

Modern sewage disposal. Closely related to the problem of main- 
taining an adequate supply of safe drinking water is the problem of 
disposing ol the domestic and industrial wastes that are produced 
especially under the crowded living conditions of modern cities. 

Sewage, as these wastes are called, includes chiefly the organic 
wastes that may be disposed of by ft nl re faction, a process that con- 
sists of a combination of bacterial action and oxidation. In some 
communities, sewage is discharged directly into streams which carry 
the sewage away. The dissolved oxygen in the water of the stream 
eventually oxidizes sewage; but, as a result, the dissolved oxygen 

Fig. 46. Water filtration. As the water passes 
through the filtering layers, solid matter and 
many germs are removed. 




An effective and relatively 
inexpensive method of puri- 
fying water is to aerate it 
by spraying it info the air. 

is used up and all the fish and other higher plant and animal life 
normally found in the stream are unable to live in the water because 
they cannot obtain oxygen. Only a few of the lower forms of life, 
such as certain algae and bacteria, can live in the polluted stream, 
and the water is unfit for almost all purposes. Oxygen-consuming 
factory wastes are sometimes handled in disposal wells. Fortunately, 
the oxygen of the air cleanses the polluted stream in the course of 
its meanderings. The distance necessary to cleanse the stream de- 
pends upon several factors the amount of sewage discharged, the 
size of the stream, its rate of flow, and so forth. 

Most modern cities dispose of sewage by more modern and less 
harmful methods of treatment such as the activated sludge process. 
The sewage, which flows to the modern treatment plant in great 
quantities of flushing water, is first run into closed tanks where the 
solids settle as a sludge, or is run through sieves and screens that 
remove the suspended solids. Certain bacteria present in the sew- 
age decompose the sludge, liberating both nitrogen and methane 
gas, and convert the sewage into a nontoxic, humus-like waste that 
may be used as a fertilizer. The partially purified water may be 
further purified by one of the methods of purification already dis- 
cussed, or it may be discharged into a neighboring stream or body of 
water. In certain types of installations, the methane gas produced 
during the digestion of the sludge is burned in gas engines to pro- 
duce the energy necessary to operate the sewage-treatment plant. 

How water can be made chemically pure by distillation, None 
of the methods mentioned removes all impurities from water. Not 
one of them completely removes substances dissolved in water. Dis- 
solved substances are completely removed from solution both in 
industry and in the laboratory by distillation. Because the boiling 
points of substances differ, it is possible to separate solute from 



inlet from 


Fig. 47. Distillation of water using a Liebig condenser. Why should 
cold water enter the jacket at the lower end? 

solvent. In general, solids that are dissolved in water have higher boil- 
ing points than water, and remain behind after the water has been 
boiled off. Certain liquids, including glycerin, have higher boiling 
points than water; others, as alcohol, have lower boiling points. 
Hence two or more liquids in solution may be separated by distilla- 
tion. Gases in solution are driven off soon after the water is heated. 

During distillation, water is first boiled, or evaporated, and 
then the steam, or water vapor, is cooled. This cooling condenses it 
into water again. Distillation is thus a double process, including both 
evaporation and condensation. The first portion of the distillate, the 
liquid that results from the condensation of the vapor, may contain 
small amounts of dissolved gases. The final portion of the distillate 
may contain small amounts of liquids or even dissolved solids whose 
boiling points are close to that of water. If these two portions are dis- 
carded, the rest of the distillate will be free of all impurities. 

Stills used for industrial purposes are made of such material as 
copper, steel, lead-lined steel, and fused silica (SiO 2 ) . 

How water is distilled in the laboratory. The laboratory appara- 
tus for water distillation consists of a flask, a Liebig condenser 
made popular in 1850 by Justus von Liebig (ton le'biK) , a German 
chemist famed for his contributions to organic and agricultural 
chemistry, and a receiver. Impure water, which, for the purpose of 
the experiment, may contain small amounts of ammonia, salt, 
and red ink, is boiled in a sidearm flask, as shown in Fig. 47. A 
thermometer indicates the boiling point of the solution. The steam 
and water vapor pass into the inner glass tube of the Liebig con- 
denser, which is surrounded by a glass jacket having a glass inlet 
and outlet for water. To condense the water vapor before it escapes 
from the inner glass tube, cold water is circulated through the outer 
tube of the condenser. As a result of this cooling, distilled water 
collects as the distillate. 

Fractional, partial vacuum, and high vacuum distillation. When 
the boiling points of the impurities are very close to the boiling 




point of the solvent, simple distillation is not very effective. For ex- 
ample, the boiling point of grain alcohol is 78C., and if separated 
from water by a single distillation, the distillate is not pure. 

Oil refiners face a difficult problem in separating the liquids that 
make up crude petroleum. The boiling points of these liquids are 
so close together that he must resort to fractional distillation. This 
consists of heating the mixture carefully and separating it into its 
various liquid fractions, which boil off at different temperatures. For 
example, the first fraction might boil off below 200C., the next be- 
tween 200C. and 220C., and so forth. In each case, it is necessary 
to purify each fraction further by additional distillations. 

It often happens that liquids cannot be separated by fractional 
distillation, because the high temperature required may char or 
otherwise decompose the solute. In such cases partial vacuum, or 
even high vacuum distillation is used. By reducing the pressure over 
the solution, its boiling point is also lowered. 

According to the kinetic theory of matter, the molecules of water 
leave the liquid and become water-vapor molecules much more 
easily when the pressure over them is decreased. Water, for exam- 
ple, under normal atmospheric pressure boils at 100C. If this pres- 
sure is lowered to one-half normal, water will boil at 82C., and 
the process of distillation can be carried through at this lower tem- 
perature. Thus, water is removed from milk in making evaporated 
milk by partial vacuum distillation. Water is removed from sugar- 
cane sap in making granulated sugar by reduced pressure distilla- 
tion. High vacuum distillation is used in making vaccines, serums, 
antibiotics, blood plasma, frozen orange juice, and for coating metals. 

The principle of distillation is used in both the huge petroleum 
"cracker" (left) and the laboratory fractionating equipment (right). 

Shell Oil Company 


Distilled water has many uses. Because distilled water is free from 
dissolved gases from the air, it is flat and insipid in taste. Its prop- 
erties are those of chemically pure water (see page 65) . Ordinarily, 
distilled water is not used for drinking purposes on land. However, 
at sea, drinking water is commonly prepared from sea water by dis- 
tillation. The economic production of drinking water from sea wa- 
ter, which contains about 3.5 percent of dissolved salts, is very old, 
dating back at least to the time of Aristotle (about 350 B.C.) . 

Today, ships of nearly all kinds obtain drinking water by distill- 
ing sea water under reduced pressure. Seagoing vessels also produce 
water for use in boilers by distilling sea water under reduced pres- 
sure. In this way, all salts are removed, and subsequent treatment re- 
moves all traces of dissolved gases. Consequently, the operation of 
the boilers is not reduced in efficiency by the formation of large 
quantities of boiler scale (see page 489) . 

Today, aviators who are forced down at sea, and all abandon-ship 
equipment of warships and merchant vessels, carry small kits that 
may be used to prepare safe drinking water from sea water easily 
and simply. Such kits produce drinking water by chemical means. 
Their chemistry is discussed on page 491. 

Distilled water is used in storage batteries, in which ordinary drink- 
ing water should not be used because of the harmful effect of its im- 
purities on the plates of the battery. Distilled water is indispensable 
in many scientific and industrial operations. In the preparation of 
vaccines, pharmaceuticals, certain dyed textiles, and C.P. chemicals, 
distilled water is commonly used. 

Heavy water, a powerful tool of research. The heavier hydrogen 
isotopes, deuterium and tritium (see page 56) , may be represented 
respectively by the symbols D and T. It has been determined that the 
nucleus of deuterium contains one proton and one neutron. There- 
fore its atomic weight is two. The nucleus of tritium contains one 
proton and two neutrons and it has an atomic weight of three. 

When water is formed from oxygen and one of the two heavy hy- 
drogen isotopes, its molecular weight is greater than 18, the molecu- 
lar weight of ordinary water. This water is called heavy water. Heavy 
water can also be formed from deuterium or tritium and any one 
of the three heavy isotopes of oxygen. Deuterium oxide, D 2 O, is 
present in ordinary water to the extent of about one part of D 2 O 
in 5000 parts of water. Tritium is seldom found in nature but is or- 
dinarily made in an atomic pile. 

Deuterium oxide has been prepared in large quantities. It differs 
to a small extent from ordinary water in both freezing and boiling 


points. Its maximum density occurs at 11.6C. rather than at 4C., 
as is the case with H 2 O. Heavy water is used in atomic piles as a 

Deuterium oxide appears to arrest the growth of seedlings; tad- 
poles die prematurely in it; it is, however, not toxic to man and the 
higher animals. 

Heavy hydrogen has been substituted for ordinary hydrogen in 
certain fats, and the course and changes which these "tagged" fat 
molecules have undergone on their way through the animal body 
have been studied by the tracer technique. 


Cerna, Wendell W. "Industrial Water Conditioning Proc- 
esses/' Journal of Chemical Education, March, 1943, pp. 107- 
115, and April, 1943, pp. 191-197. 

Ellis, Cecil B. Fresh Water from the Ocean. Ronald Press 
Company, New York, 1954. This is a conservation study deal- 
ing with water for cities, industry, and irrigation. 

Goldblatt, L. A., Ed. Collateral Readings in Inorganic Chem- 
istry. D. Appleton-Century Co., New York, 1937 (2nd series, 
1942). No. 8 of the 31 articles in this collection deals with 
"Factors Contributing to Quality of Public Water Supplies." 
Written by H. E. Jordan. 


1. A solution is a uniform mixture of solvent and solute 
which does not conform to the law of definite proportions. 

2. A dilute solution contains very little solute in comparison 
with the solvent; a concentrated solution contains a large 
amount of solute. 

3. Some of the factors which determine the amount of solute 
that will dissolve in a solvent are: (1) the nature of solute 
and solvent, (2) temperature, and (3) pressure. The speed of 
solution depends upon (1) the state of subdivision of the 
solute, and (2) how thoroughly the solute and solvent are 

4. All solutes raise the boiling point and lower the freezing 
point of the solvent. The amount of solute determines the 
number of degrees of change. Acids, bases, and salts in solution 
affect boiling and freezing points to a greater degree than do 
other substances, including alcohol, sugar, and glycerin. 


5. Solutions and suspensions differ in the following respects: 

(1) In a suspension, the mixture separates on long standing. 

(2) The particles in suspension are much larger than the 
particles of a solute in a solvent (which are in the molecular 
state of subdivision) . (3) The particles of a suspension may 
be separated out by filtration; the particles of a solution can- 
not be so separated. (4) A solute has greater effect on the 
freezing and boiling points than has a material in suspension. 

6. The colloidal state of matter is in the borderland be- 
tween true solutions and suspensions. 

7. Distillation consists of evaporation and condensation. 
Impurities can be removed by distillation because of the dif- 
ference in the boiling points of a solvent and its solute. 


Group A 

1. (a) What is the purest form of H 2 O found in nature? 
(b) What property of H 2 O makes it almost impossible to find 
pure H 2 O in nature? 

2. (a) What is a solution? (b) In a solution of NaCl in 
H 2 O, which is the solvent and which the solute? (c) What is 
a tincture? 

3. Distinguish a dilute solution from a concentrated so- 

4. Name four factors that determine the degree and the 
speed of solubility of a substance. 

5. Why is more sugar or more stirring required to sweeten 
iced coffee than hot coffee? 

6. Why should sealed bottles of H 2 O heavily charged with 
CO 2 be kept cold? 


7. How would you proceed to prepare a saturated water 
solution of washing soda, Na 2 CO 3 ? 

8. How could you change a saturated solution of CuSO 4 to 
an unsaturated solution having the same volume? 

9. Without changing the temperature or volume of a solu- 
tion of CO 2 in H 2 O, how could you change an unsaturated 
solution of this gas into a saturated one? 

10. How could you determine whether a solution is satu- 
rated, unsaturated, or supersaturated? 

11. Mention one commercial use of supersaturation. 


12. What is the effect of a solute on the freezing and boiling 
points of a solvent? 

13. What principle is involved in the use of an antifreeze 
mixture in an automobile radiator? 

14. How do true solutions and suspensions differ? 

15. What is the name given to the state of matter that is on 
the borderland between true solutions and suspensions? 

16. What are seven methods used to purify drinking water? 

17. How is sewage disposed of in modern sewage-disposal 

18. What may be the results of ineffective methods of sewage 

. . I . . . 


19. Which method or methods of purifying H 2 O produce 
chemically pure H 2 O? 

20. (a) What is distillation? (b) Upon what fact does the 
separation of impurities by distillation depend? 

21. Why is the middle portion of a distillate the purest? 

22. Make a diagram of the apparatus used in the laboratory 
distillation of H 2 O. 

23. (a) Is the drinking water of a large city, such as Chicago 
or New York, distilled? (b) Give reasons for your answer. 

24. Petroleum is refined by fractional distillation. Why? 

25. Why are liquids often distilled under reduced pressure? 

26. State the physical properties of distilled water. 

27. For what is distilled water used? 

28. What are the desirable characteristics of drinking water? 

29. What is the difference between wholesome water and 
chemically pure water! 

30. (a) What is heavy water? (b) How does it differ from 
ordinary water? (c) Name two uses of heavy water. 

Group B 

31. Devise an experiment to show that perfectly clear spring 
water contains impurities. 

32. Small filters attached to household faucets sometimes 
become a menace to health. Explain this statement. 


33. What would be the effect of the continued use of rain 
water in storage batteries? 

34. Is water obtained by melting ice from a lake purer than 
the water of that lake? Explain your answer. 

35. If a liquid is colorless, odorless, and clear, how could 
you determine whether it is a solution or a pure compound? 

36. How does a solution of NaCl in water differ from a mix- 
ture of NaCl and sugar? 

37. Can a dilute solution be a saturated solution also? Ex- 

38. Explain the operation and principle of the pressure 

39. (a) Explain solution by means of the kinetic theory of 
matter, (b) According to the kinetic theory, explain why it is 
easier to evaporate or distill a liquid under reduced pressure. 

40. Would there be any advantage or disadvantage in 
making solutions of (a) table salt, (b) calcium hydroxide, 
with hot water instead of cold? (c) Explain each answer. 


1. Visit your local sewage-disposal plant, and study its op- 
eration. Make a report on this visit to your class. 

2. Obtain some Rochelle salt (U.S.P.) from your teacher 
or druggist and prepare pure crystals from it. Bring them to 
class and explain the process you used. (As a substitute, you 
may prepare "rock candy" crystals.) 

3. Harold C. Urey, discoverer of heavy water, in a lecture in 
1938, made the following statement, "I believe I speak for the 
vast majority of all scientific men. Our object is not to make 
jobs and dividends. These are a means to an end, merely inci- 
dental. We wish to abolish drudgery, discomfort and want 
from the lives of men, and bring them pleasure, comiort, 
leisure and beauty. Often we are thwarted but in the end we 
will succeed." Write a short report either for or against this 
view, or organize a class discussion on this subject. 

4. Study the methods of water purification in your own 
community. Report to your class on this subject. Find out, if 
you can, the mineral content of the water and the amount and 
kind of suspended solids present both before and after treat- 

5. Using hypo (sodium thiosulfate) prepare 10 ml. of a 
supersaturated solution. Allow the solution to cool slowly to 
room temperature. Then add a crystal of hypo. Repeat this 
before your class if it can be arranged. 



. . . / heard Cleve say: "Do you 
believe sodium chloride is dissolved 
into sodium and chlorine? In this 
glass I have a solution of sodium 
chloride. Do you believe there are 
sodium and chlorine in it?" "Oh, 
yes," Ostwald replied, "there is some 
truth in that idea." . . . Cleve threw 
a look at Ostwald which clearly 
showed that he did not think much 
of his knowledge of chemistry. 

Svante Arrhenius, 1925 

Two eternal questions: How? and Why? Science is constantly try- 
ing to answer two questions how and why. Often it is not too dif- 
ficult to answer the hows but the whys well, that is a different 
story. Theories must be formulated, tested, and adapted to keep 
them in accord with all the observed facts, and that is a big job. 

At least two great theories underlie much of chemistry: the elec- 
tron theory and the theory of dissociation. In a sense, the theory of 
dissociation is but an aspect of the electron theory, for the theory 
of dissociation is explained in terms of the electron theory. These 
two great theories explain some of the hows and whys of chemistry. 

A Swedish boy tackles some puzzling questions. Why is distilled 
water a nonconductor of electricity? Why do water solutions of some 
substances conduct electricity yet water solutions of other substances 
do not? What causes some acids and bases to be strong while others 
are weak? 

These were some of the unsolved problems that confronted 
chemists when Svante Arrhenius (ar-ra'ni-us) was still at school in 

Arrhenius not only wondered about these problems but set to work 
to solve them. He had some unusual notions of his own about the 



way in which electricity passes through solutions. Day after day, and 
often far into the night, he worked in his laboratory with hundreds 
of different solutions. For two years he toiled ceaselessly. 

Arrhenius attempted to formulate a theory that would explain 
what he had observed. In those days his whole world, both of wak- 
ing and sleeping hours, was one of solutions, electric currents, and 
formulas. For him the rest of the world did not exist. One night he 
sat up very late. Suddenly, like a flash, he saw the answer to the great 
riddle. "I got the idea/' he wrote, "in the night of the 17th of May 
in the year 1883, and I could not sleep that night until I had worked 
through the whole problem." 

He went to his teacher of chemistry. "I have a new theory of 
electrical conductivity/' Arrhenius told him. The professor looked 
at this boy and said: "You have a new theory? That is very interest- 
ing. Good-by." But Arrhenius did not lose heart. He wrote to the 
leaders in chemistry. Most of them were hostile to his revolutionary 
theory. After a long struggle, however, they were forced to admit its 
probable truth and later saw Arrhenius awarded the Nobel prize in 
chemistry. And so it happened that a mere boy, with the clear insight 
and the creative imagination of a truly great scientist, stepped in and 
cleared away an obstacle that had stood squarely in the path of chem- 
ical progress. 

How Arrhenius explained the conductivity of solutions. The the- 
ory proposed by Arrhenius is known as the theory of ionization, or 
dissociation. Arrhenius assumed that when an electrolyte, such as 
sodium chloride, dissolves in water, it tends to dissociate, or ionize. 
That is, it tends to break apart into electrically charged atoms or 
groups of atoms (radicals) . Arrhenius used the term ion, meaning 
wanderer, to refer to an atom or group of atoms carrying an electric 
charge. He represented the dissociation of the sodium chloride mole- 
cule into ions when it dissolves in water thus: 

NaCl Na+ (sodium ion) + Cl~ (chloride ion) 
He represented the dissociation of the sodium nitrate molecule thus: 

NaN0 8 ->Na + +(NO 3 )- 

Arrhenius could not see the ions in solution. They are far too small 
to be seen. He advanced the theory that they were present because 
he could account for what he observed only by assuming such ions to 
be present. To him, when an electrolyte dissolves, a certain number 
of its molecules immediately split up into ions. Thus an electrically 


neutral compound tends to dissociate into a number of positively and 
negatively charged particles, or ions. They move about in all direc- 
tions until the passage of an electric current draws each ion to the 
electrode bearing an opposite charge. 

It is the ions that carry an electric current, or flow of electrons. 
Hence a substance that does not dissociate into ions is a nonelectro- 
lyte, and a substance of which a large proportion dissociates in water 
solution is a good electrolyte. Arrhenius' theory of ions and of elec- 
trolytic dissociation is today, with certain modifications, universally 
accepted as the correct explanation of the conductivity of solutions. 

How an ion differs from an atom. There are two main differences 
between an atom and an ion: (1) An atom is electrically neutral; an 
ion is positively or negatively charged. (2) An atom always consists 
of a single element; an ion may consist of one or more than one ele- 
ment, as in the case of the ammonium ion, (NH 4 ) +. A sodium atom 
is quite different from a sodium ion. The former is a silvery metal 
particle that reacts violently with water; the latter is a colorless par- 
ticle that has no noticeable reaction with water. 

The dissociation of zinc chloride may be represented thus: 

Note the change from C1 2 in ZnCl 2 to 2 Cl~ when the zinc chloride 
is dissociated into ions. The number of positive or negative charges 
on the ion is equal to the valence of the element or radical. 

showing what happens when the following electrolytes dis- 

a) Nitric acid, HNO r g) Sulfuric acid, H 2 SO 4 . 

b) Lithium hydroxide, LiOH. h) Potassium carbonate, 

c) Potassium hydroxide, K 2 CO 3 . 

KOH. i) Sodium phosphate, 

d) Ammonium hydroxide, Na 3 PO 4 . 

NH 4 OH. j) Potassium chlorate, KC1O 3 . 

e) Barium hydroxide, k) Magnesium bicarbonate, 
Ba(OH) 2 . Mg(HC0 8 ) 2 . 

f) Aluminum chloride, A1C1 3 . 1) Calcium nitrate, 

Ca(N0 3 ) 2 . 

Where and how are free ions formed? Most commonly, dissocia- 
tion takes place in a water solution. However, it also occurs in com- 
pounds such as sodium hydroxide, NaOH, when they are heated un- 
til they melt, or fuse. Such compounds are made up, not of molecules, 
as was previously supposed, but of ions in a definite pattern (see 


direction of flow 
of electrons 

heated filam 

Fig. 48. Simplified diagram of a radio 
tube showing the hot filament emitting 
a stream of electrons. 

page 631). As the compound is heated, its ions vibrate more rap- 
idly and eventually tear apart, leaving them free to carry an electric 
current. Gases, too, at very high temperatures dissociate into ions, 
thus becoming conductors of electricity. Flames, x-rays, and radio- 
active elements ionize the air around them. 

A metal, when heated to incandescence, boils off electrons from 
the outer rings of its atoms. This principle is used in the electron 
tubes so familiar to us in radios, television sets, and other electronic 
instruments. Certain types of electron tubes are filled with a gas or 
mixture of gases at low pressure. Electrons from the hot cathode of 
the tube strike the molecules of the gas, ionizing them. 

The positive ions move toward the negative electrode, or cath- 
ode, and the negative ions and electrons move toward the positive 
electrode, or anode. In a radio tube, the anode is the plate, and 
the cathode may be the hot filament. Thus a stream of electrons or 
electric current flows through the tube. Because a gas-filled tube 
contains more molecules that can be ionized than does a vacuum 
tube from which most of the molecules have been removed, relatively 
large currents can flow through a gas-filled tube. The thyratron tube, 
often employed as a rectifier, is a common example of a gas-filled 
electron tube. 

How Arrhenius explained the action of acids. All acids in water 
solution contain free hydrogen ions. These hydrogen ions determine 
the typical properties of acids. It is the free hydrogen ion that turns 
blue litmus pink, has a sour taste, is replaceable by a metal, and 
neutralizes a base by combining with its hydroxyl ion. 

The number of hydrogen ions present determines the strength of 
an acid. A strong acid is one that dissociates easily, and thus pro- 
duces a large number of hydrogen ions. The most commonly used 
strong acids are hydrochloric, nitric, and sulfuric acids. Carbonic 
acid, H 2 CO 3 , sulfurous acid, H 2 SO 3 , and boric acid, H 3 BO 8 , are weak 
acids, because they do not dissociate easily, and thus form only a 
small number of hydrogen ions in solution. 


How Arrhenius explained the action of bases. What has been 
said about strong and weak acids refers equally well to strong and 
weak bases, except that the determining factor here is the number of 
free hydroxyl (OH) ~ ions present. In water solution, a strong base, 
such as sodium hydroxide, NaOH, potassium hydroxide, KOH, or 
lithium hydroxide, LiOH, forms large numbers of hydroxyl ions. 
A weak base, such as ammonium hydroxide, NH 4 OH, produces a 
comparatively small number of (OH)- ions. 

Some bases, like copper hydroxide, Cu (OH) ,, and aluminum hy- 
droxide, Al (OH) 3 , are extremely weak, because, in addition to the 
fact that they form a relatively small number of ions, they are also 
only sparingly soluble. The properties of bases are caused by the pres- 
ence of free hydroxyl ions. Hence dry, solid sodium hydroxide has no 
basic properties; that is, it does not react as a base. 

How the theory of Arrhenius explained neutralization. The hy- 
drogen ion and the hydroxyl ion are responsible for acidic and basic 
properties respectively. If these free ions are removed, acidic and 
basic properties are destroyed. This is exactly what happens in neu- 
tralization. The hydrogen ion of the acid unites with the hydroxyl 
ion of the base, forming water, ;i nonelectrolyte. 

Na+ + (OH)~ + H+ + Cl~ -> Na+ + Cl~ + HOH 

Hydrogen-ion concentration (pH) and how it is measured. It is 

frequently important to know whether complete neutralization has 
been produced, or what the degree of acidity or alkalinity (basic- 
ity) of a substance is. Sugar-refiners, brewers, paper-makers, electro- 
platers, sanitary engineers, and bacteriologists must have a working 
knowledge of the acidity or alkalinity of many substances. 

Chemists make use of a hydrogen-ion scale in which the unit is 
the /;H value (pronounced by reading the letters p, H) , just as the 
unit of temperature is a degree. A jm value of seven is considered 

Chemists in professional labora- 
tories use special equipment to 
determine the pH value of sol- 



true neutralization. Pure water has such a value. A pn value less than 
seven indicates an acid condition. Thus, saliva has a pn value of 
about 6.9 and is slightly acid. A pn value greater than seven indi- 
cates an alkaline condition. Thus, normal blood has a pu value of 
about 7.3 and is slightly alkaline. A value higher than 7.5 repre- 
sents a condition of alkalosis. Since the human blood is normally 
slightly alkaline, a pn below 7.2 indicates a condition of acidosis. 
A pn value of five represents an acidity 10 times as great as a pn 
of six. 

Litmus may be used as an indicator only with those solutions 
whose pn is not less than four or greater than eight. Other indi- 
cators have wider or narrower ranges. By comparison with standard 
color tubes for each indicator, the degree of acidity or alkalinity 
of a liquid can be determined accurately. 

Several types of indicator papers have been developed for deter- 
mining the pn values of solutions. The use of such papers contain- 
ing water-soluble dyes is very convenient and is much quicker than 
referring to standard color tubes. One such paper is known as 
"Hydrion Paper." 

How the degree of dissociation was determined. We have seen 
that some substances produce a large percentage of ions and other 
substances produce a small percentage of ions. In other words, com- 
pounds differ in the extent to which they form ions in solution, that 
is, in their degree of dissociation. The degree of dissociation of a 
solution depends upon (7) the solute, (2) the solvent, (3) the con- 
centration of the solution, and (4) the temperature. 

The degree of dissociation is measured by electric conductivity, 
that is, the ease with which an electric current passes through a 
solution. This may be determined roughly by setting up an appa- 
ratus as shown in the illustration below (Fig. 49) . When the bottle 
contains carbonic acid (a weak acid) the electric-light bulb glows 
faintly. When it contains a dilute solution of hydrochloric acid (a 
strong acid) , the light bulb glows brightly, showing that little resist- 
ance is being offered to the passage of the current. A solution of 
sugar in water, when placed in the bottle, does not produce a glow 
in the bulb, and thus shows that a sugar solution is a nonelectrolyte. 

Fig. 49. Laboratory setup for studying 
the conductivity of a liquid. When will 
the bulb light? 

platinum electrodes 
liquid to be tested 


Arrhenius determined the apparent percentage of dissociation of 
many compounds by this method. He found that in one-tenth nor- 
mal water solution (0.1N) the apparent dissociation of hydrochloric 
acid, nitric acid, and sodium hydroxide was about 92 percent (strong 
electrolytes) ; that of sulfuric acid about 61 percent; that of potas- 
sium chloride about 86 percent; and that of acetic acid, ammonium 
hydroxide, and mercuric chloride only about 1 percent (weak elec- 
trolytes) . Later it was found that this method had certain limitations. 

How Arrhenius explained abnormal boiling and freezing points 
of solutions. The theory of ionization also explains why certain 
ac ids, bases, and salts raise the boiling points of their water solu- 
tions to an abnormally high degree, though sugar does not to the 
same extent. The acid dissociates and produces two or three times 
as many particles (ions) as there are molecules of undissociated sugar. 
The higher percentage of dissociation of the acid produces a greater 
increase in the boiling point, since it is the actual number of par- 
ticles (ions or molecules) in solution that determines both the boil- 
ing and freezing points of a solution. 

How the theory of ionization fits in with the electron theory. 
When Arrhenius proposed his theory of dissociation, the electron was 
still undiscovered, and the electron theory of matter had not yet 
been formulated. Ions, however, fit in beautifully with our present 
electron concept of matter. Atoms become ions by gaining or losing 
electrons. For example, if an electrically neutral atom of sodium 
(Na) loses 1 electron, it becomes a positively charged particle, and 
we represent it Na+. This is a sodium ion (Na > Na+) . Sim- 
ilarly, if an electrically neutral atom of chlorine (Cldb) gains I elec- 
tron, it becomes a negatively charged particle, and we write it Cl~. 
This is a chloride ion (Cl + -> Cl~) . 

The compound made up of these elements (NaCl) is really com- 
posed of both of these ions held strongly together, making what is 
known as an ion-molecule that is electrically neutral. However, when 
these ion-molecules are placed in water, they are split into two parts, 
and so long as water is present, we have separate sodium ions and 
chloride ions. Dissociation, then, may be represented as shown below. 

Fig. 50. Dissociation of a sodium chloride molecule, ac- 
cording to the electron thepry. 

NaC1 -molecule *- Na+ (ion) + CT (ion) 

(+11-11-0) (-H7-17-0) (+11-10-+1) (+17-18--1) 


The charge on the ion is thus seen to be the same as the valence of 
the element. 

How the theory of Arrhenius was later modified. In general, 
Arrhenius' theory has stood the test of time very well. In three re- 
spects, however, it has been modified slightly. 

1) The Swedish chemist thought that water simply kept the ions 
apart, but today we have a better understanding of the impor- 
tant role of the solvent. The covalent water molecule is pictured 
as a tiny magnet with a plus hydrogen end and a less positive oxy- 
gen end. Such a molecule is called a dipole. 

When crystals of the polar compound, NaCl (made up of a 
lattice of Na and Cl ions see page 630) , are added to water, the 
positive end of the dipole molecule attracts the negative Cl ion; 
the other end of the dipole attracts the positive Na ion. The Na 
and Cl are thus dissociated and free sodium and chloride ions are 
produced in solution. When the water is evaporated, the ions re- 
combine forming ion-molecules of NaCl crystals. NaCl and other 
ionic compounds are strong electrolytes. 

2) The Br0nsted-Lowry theory has added to what Arrhenius be- 
lieved about the ionization of acids. According to this theory, an 
acid is a proton (hydrogen ion) donor; that is, any compound 
that tends to lose a proton (H+) to another substance is an acid. 

For example, in the case of a water solution of HC1, the water 
molecule combines with the positive hydrogen ion end (proton) 
of the HC1, forming a hydrogen (hydronium) ion, leaving the 
chlorine negatively charged, thus: 

H 2 O + H Cl ^H 2 . H+ + Cl" 

V^^CI) hydronium ion 

As HC1 dissolves in water, heat is liberated, indicating a chemical 
change is taking place. The hydronium ion is an acid since it may 
give up its proton. It is also written as H 8 O+. 

A base is a proton acceptor. A base is any substance which com- 
bines with a hydrogen ion or proton. Thus the water molecule 
which will accept a proton is a base, though not an active one. 
Bases are often negative ions such as OH~, which is an active one. 

3) Insofar as weak electrolytes are concerned, the theory ex- 
pressed by Arrhenius is still correct. They are not completely dis- 
sociated. In the case of strong electrolytes, however, it is now be- 
lieved that they dissociate completely. The fact that the heat of 

4(OH)-*-2H 2 

hydrogen^ ^JB^Wiiii^^Xoxygen 

Fig. 51. Action during the 
electrolysis of water. 

neutralization of all strong acids and bases is the same is one evi- 
dence in support of this belief. The fact that when dilute solu- 
tions of strong electrolytes are mixed there is no trace of heat is 
further proof. The degree of dissociation of weak electrolytes can 
be determined accurately, but that of strong electrolytes cannot 
because of the disturbing electrical effects of a large concentration 
of ions. In such a high concentration, ions do not behave as "free 
ions" and only an apparent ionization is obtained. 
The theory of dissociation as stated by Arrhenius is still useful. 
Other new ideas have added to the accuracy of our concept of dis- 
sociation. For example, we know that molecules of compounds such 
as carbon dioxide do not dissociate in solution and are nonelectro- 
lytes because they are nonpolar. However, these new ideas are of 
value chiefly in dealing with phenomena which are beyond the scope 
of this book. For this reason, we shall follow the theory of ioniza- 
tion as Arrhenius originally presented it. 

What happens during the electrolysis of water? So far we have 
considered pure water a nonelectrolyte. Roughly speaking, this is 
true, since in about 1 billion grams of water only about 1 gram of 
free hydrogen ions is present. Yet this very slight dissociation of wa- 
ter [H,O H+ + (OH) -] is important. Water is such a poor con- 
ductor that acid must be added to conduct the current. Nevertheless, 
the hydrogen and oxygen that are products of electrolysis come from 
the water. How? 

The positive hydrogen ions travel to the negative cathode. The 
cathode, connected to a battery, is supplied with electrons. When the 
hydrogen ion reaches the cathode, an electron, e, from the cathode 
is given up to the hydrogen ion, which changes to a neutral hydro- 
gen atom. This atom immediately joins another hydrogen atom and 
is liberated as a molecule of gas, thus: 

H+ + c - H; H + H - H 2 T 



The negative sulfate ions, and hydroxyl ions from the water, travel 
to the positive anode. The hydroxyl ion gives up its electron more 
readily than does the sulfate ion. Having lost its electron, it becomes 
the hydroxyl radical, which breaks down, forming water and oxygen. 
The sulfate ions remain in solution. The concentration of the acid 
remains unchanged. 

(OH)---(OH); 4(OH)->2H 2 + 2 ! 

What is hydrolysis? We should expect neutral salts to show 
neither acid nor basic properties, since they contain neither hydrogen 
nor the hydroxyl radical which might form ions. Yet when we add 
blue litmus to a solution of copper sulfate, the litmus turns pink, in- 
dicating an acid. Where are the free hydrogen ions to account for 
this behavior? The answer lies in the fact that water is slightly dis- 

Cu++ + (SO 4 )~ + 2HOH = 2H+ + (SO 4 )~ + Cu(OH) 2 

Since the copper hydroxide formed is only very slightly dissoci- 
ated, there will be some union of Cu++ and (OH) -, thus forming this 
very weak base, and liberating an excess of hydrogen ions that form 
the strong acid, sulfuric acid. Therefore, the solution* is slightly acid. 
This reaction is the reverse of the neutralization of sulfuric acid by 
the base Cu (OH) ,. Hydrolysis of a salt is the action between the ions 
of water and the ions of a dissolved salt, forming an acid and a base. 

Hydrolysis also explains the basic nature of Na 2 CO 3 , thus: 

2Na+ + (C0*)~ + 2HOH <= 2Na+ + 2(OH)~ + H 2 CO 8 

Since the H 2 CO 3 is a weak acid, there will be some union of H+ 
and (CO 8 ) , thus forming the slightly dissociated H 2 CO 3 , and lib- 
erating an excess of OH ions which form the strong base, NaOH. 

A solution of sodium chloride and water is perfectly neutral, since 
it is made from a strong acid and a strong base. A solution of a salt 
made from a strong base and a weak acid, or from a weak base and 
a strong acid, does not show a neutral reaction. 


Complete the following equations, and tell whether the 
solution in each case will be acidic, basic, or neutral toward 


a) K 2 CO 3 + 2HOH - d) Na 2 SO 4 + 2HOH - 

b) ZnSO 4 + 2HOH- e) KNO 3 + HOH -> 

c) A1C1 3 + 3HOH -> f) K 2 SO 4 + 2HOH - 

What are reversible reactions? Most reactions we deal with are 
reactions that take place in water solutions and, hence, are reactions 
between ions. The presence of free ions facilitates chemical changes. 
If we dissolve sodium chloride, NaCl, and potassium nitrate, KNO 3 , 
in water, these salts immediately dissociate, forming free Na+, Cl~, 
K+, and (NO 3 ) ~. These swiftly moving ions constantly meet and 
form molecules of KNO 3 and NaCl. The entire reaction, therefore, 
is said to be reversible, that is, it goes in both directions. 

KNO 3 + NaCl * KC1 + NaNO 3 

A reversible reaction always reaches a point at which change is no 
longer apparent. In other words, the reaction has reached a point of 
balance or equilibrium. This does not mean that nothing is happen- 
ing. On the contrary, the equilibrium is dynamic, or moving, for the 
substances are breaking up as rapidly as they are being formed. 

Reversible reactions and equilibrium. The dissociation of an elec- 
trolyte is a reversible reaction. Thus, when acetic acid, HC 2 H 3 O 2 , is 
dissolved in water, free hydrogen ions and free acetate ions are 
formed. These ions meet and form acetic acid. 

HC2H 3 2 *= H+ + (CzHsOz)- 

Finally a time is reached when the rate of change from free ions 
to HC 2 H 3 O 2 will just equal the rate of change from HC 2 H 3 O 2 to 
free ions. This will happen, in the case of a 0.0 IN acid solution, 
when 98.7 percent of the HC 2 H 3 O 2 is in the form of free ions and 1.3 
percent is in the form of HC 2 H 3 O 2 . When this condition is reached, it 
would appear that the HC 2 H 3 O 2 is no longer dissociating and that a 
condition of stable equilibrium has been reached. In reality the equi- 
librium is not stable but dynamic, or changing. Changes go on even 
after the 98.7 : 1.3 ratio is attained, but while new HC 2 H 3 O 2 is being 
formed, more is being ionized, keeping the same ratio. 

When do reactions go to completion? Substances that do not 
dissociate cannot react reaoily with other substances that do. If even 
one of the products of a reaction is unable to dissociate to any great 
extent, the backward action cannot take place. The reaction is then 
said to go to completion. Advantage is taken of this to secure reac- 
tions that complete themselves. 

Fig. 52. Action of silver nitrate and dilute hydrochloric acid. 
Does this reaction go to completion? Why? Under what 
conditions do reactions go to completion? 

Insoluble substances, called precipitates, gases liberated under nor- 
mal temperature conditions, and pure water are practically incapable 
of dissociation. Hence a reaction will go to completion whenever one 
of the products formed is (1) a precipitate, (2) a gas, or (2) water. 
Examples of such reactions follow: 

1) Formation of a precipitate. Chemists use as tests reactions in 
which precipitates are formed. Thus, in testing for a chloride, silver 
nitrate, AgNO ;{ , is added. The formation of insoluble silver chloride, 
AgCl, prevents the reaction from reversing itself. 

Ag+ + (N0 3 )- + Na+ 

AgCl [ + Na+ + (NO 8 )- 

2) Formation of a gas. IrTthe laboratory preparation of hydrogen 
chloride, a mixture of sodium chloride and sulfuric acid is heated, 
and the hydrogen chloride gas that is liberated leaves the field of ac- 
tion. Hence the reaction does not reverse itself. 


l- + H+ + (HS0 4 ) ~ - HC1 

(HSO 4 ) 

3) Formation of undissociated water. During neutralization, wa- 
ter is always one of the products formed. Therefore the neutralizing 
reaction is complete, since water, which is practically undissociated, 
may be considered as having left the field of chemical action. 

Na+ + (OH)- + H+ -f Cl- - Na+ -f Cl~ + HOH 
This is strictly true only when the salt formed is not hydrolyzed. 


1. Complete and balance the following equations: 

a) KNO 3 + NaCl - c) NaOH + PbCl 2 -> 

b) BaCl 2 -f K 2 SO 4 - d) KOH 4- H 2 SO 4 - 



e) Na 2 S0 3 + H 2 S0 4 -> SO, + H 2 O + 

f) FeS 4- HC1 -> 

g) CaCO 3 -f HN0 3 -> CO 2 + H 2 O + 
h) CaCl 2 + NaN0 3 -> 

i) Na 2 SiO 3 -f Ca(OH) 2 -> 

2. Examine the foregoing equations and, in each case, see 
whether any of the products formed are gases or precipitates. 
Mark them with the appropriate arrows, j or j. Remember 
that S0 2 , H 2 S, and CO, are gases. 

3. Finally, with the aid of the following table of solubilities, 
tell whether each of the reactions goes to completion or not, 
stating your reason in each case. 

TABLE 10. 

All nitrates, chlorates, and acetates are soluble in water. SOLUBILITY OF 

All chlorides, bromides, and iodides are soluble, except those of Ag, Pb, and Hg. JULUBILI T ur 

All sul/ates are soluble, except PbS0 4 , BaS0 4 , CaSO,, HgSO,, and Ag,S0 4 . COMMON 

All hydroxides are insoluble, except those of Na, K, NH,, Ca, and Ba. COMPOUNDS 

All sulfides are insoluble, except those of Na, K, NH 4 , Ca, and Ba. 
All oxides are insoluble, except those of Na, K, and Ca. 
All carbonates, silicates, and phosphates are insoluble, except those of Na, K, and NH.,. 

The law of mass action and equilibrium, The quantity of a sub- 
stance in a unit volume of solution is a measure of the concentration 
of a solution. As early as 1803, the French chemist Claude Berthollet 
(who first made use of the bleaching action of chlorine) noticed that 
the direction of a chemical reaction is dependent upon the concentra- 
tions of the substances involved in the reaction. 

He noticed that, in general, the greater the concentration, the 
greater the speed of the reaction. For example, a match burns quietly 
in ordinary air which contains about 21 percent oxygen. In pure oxy- 
gen, however, the match burns much more quickly, since the con- 
centration of the oxygen (one of the reacting substances) has been 
increased almost fivefold. A greater concentration of oxygen means 
more molecules of oxygen per unit volume of gas, and, therefore, a 
greater possibility for oxygen molecules to come in contact with 
molecules of carbon and carbon compounds. This causes an increase 
in the speed of the chemical reaction involved in burning. 

Sixty-four years later, Guldberg and Waage, professors of mathe- 
matics and chemistry at the University of Oslo, Norway, demon- 
strated that the speed of a chemical reaction is directly proportional 
to the concentrations of the reacting substances. They also made an 
interesting discovery concerning the point of dynamic equilibrium 



of a chemical reaction that is, the poiixf at which the reaction 
in one direction just balances the reaction in the opposite direction. 
They found that a chemical reaction which is normally reversible 
can be forced to go in one direction with small reversal. This is ac- 
complished by manipulating the concentrations of the reacting sub- 

Guldberg and Waage expressed this phenomenon in the form of 
the law of mass action, which implies that a change in the quantity 
of the reacting substances results in a change in the equilibrium point 
of the reaction. In the manufacture of chemicals, the direction of 
a reaction is so controlled that large yields are produced. 

How the addition of a common ion forces a reaction to go to 
completion. In the light of the law of mass action, let us consider 
a saturated solution of sodium chloride. We may express the reac- 
tion that is taking place as: 

NaCl <=> Na+ + Cl~ 

In such a reaction, the product of the concentrations of the free so- 
dium ions and the free chloride ions is a constant. If, by any means, 
we increase the number of chloride ions, the number of sodium ions 
must decrease. The number of sodium ions can decrease only if some 
of the sodium chloride comes out of solution. 

To increase the chloride ions, we add to the solution a compound 
of chlorine that dissociates to a high degree. Hydrochloric acid is 
such a compound. Therefore, if hydrochloric acid is added to a 
saturated solution of sodium chloride, some of the sodium chloride 
will be precipitated. We call this shifting of the equilibrium point 
the common-ion effect. In this case, the common-ion effect is caused 
by the addition of the chloride ion, which is common to HCi and 

Fig. 53. Apparatus used to 
show the common-ion effect. 

NaCl + H 2 SO 4 

J^ saturated 
NaCl solution 

precipitated out 

Svante Arrhenius (right) and his close friend, Wilhelm Ostwald. Each was awarded 
a Nobel prize for chemistry, Arrhenius in 1903, Ostwald in 1909. 

Almost insoluble salts, such as silver chloride, AgCl, may be com- 
pletely precipitated by adding a common ion. Thus, the addition of 
an excess of NaCl increases the concentration of the few dissociated 
chloride ions from the silver chloride, and causes some more AgCl 
to precipitate. 

A weak acid may be weakened by adding a salt of the weak acid. 
Thus the addition of sodium citrate to citric acid weakens that acid, 
because the addition of the common ion (the citrate ion) forces 
more of the citric acid to the undissociated form. 


Jaffe, Bernard. Crucibles: The Story of Chemistry, pp. 219- 
211. Simon and Schuster, New York, 1948. The story oi the life 
and work of Arrhenius. 

Samrnis, Constance S. "How Annabella Learned the Facts 
about pH." Journal of Chemical Education, Oct., 1942, pp. 490- 
494. A delightful, cleverly illustrated treatment of a very im- 
portant topic. 


1. Science must always answer at least two questions 
how? and why? 

2. Two great theories underlie much of chemistry: (1) the 
electron theory and (2) the theory of dissociation. 



3. The theory of ionization was advanced to explain elec- 
trolysis, nonelectrolytes, strong and weak acids and bases, and 
other puzzling facts. 

4. An ion is an atom or radical that carries an electric 

5. When an acid, base, or salt goes into solution, it dis- 
sociates, partially forming free ions. When a current is passed 
through such a solution, the ions are attracted to the electrodes 
having a charge opposite their own. 

6. Free ions are formed in greatest numbers in water so- 
lutions of electrolytes. Free ions occur also in molten sub- 
stances, in heated gases, and in the air surrounding radio- 
active substances. Atoms of glowing metals throw off electrons; 
if the electrons come in contact with a gas, as in a radio tube, 
the molecules of the gas become ionized. 

7. In terms of ionization, an acid is a substance that pro- 
duces free hydrogen ions; a base is a substance that produces 
free hydroxyl ions. 

8. A strong acid is one that dissociates easily, forming large 
numbers of hydrogen ions. A weak acid forms only a small 
number of hydrogen ions. Strong bases form a large number of 
hydroxyl ions; weak bases form few. A compound, such as 
water, whose molecule acts like a tiny magnet, is called a 
dipole. The molecules of nonpolar compounds, such as CO 2 , 
do not dissociate in solution, are nonelectrolytes and noncon- 
ductors of electricity. 

9. Neutralization is the union of the hydrogen ions of an 
acid with the hydroxyl ions of a base. 

10. The acidity or alkalinity of a solution is measured in pH 
values. Water, with a pH of seven, represents neutrality. A sub- 
stance with a pn greater than seven is basic; one with a pn less 
than seven is acidic. 

11. Substances dissociate to different degrees, depending 
on the (1) solute, (2) solvent, (3) concentration of the solu- 
tion, and (4) temperature. 

12. All water solutions of acids, bases, and salts are elec- 
trolytes; that is, they dissociate and conduct an electric current. 

13. Electrolytes raise the boiling points and lower the 
freezing points of solutions to a greater degree than non- 
electrolytes do, because they dissociate, forming larger num- 
bers of ions. 

14. The electron theory upholds the theory of dissociation. 
The charge on an ion is the same as the valence of the element. 

15. Hydrolysis of a salt is the action between the ions of 
water and the ions of a dissolved salt, forming an acid and a 
base. It is the opposite of neutralization. Salts formed by the 


reaction of a strong acid and a weak base are acidic; by the 
reaction of a weak acid and a strong base are basic. 

16. A reversible reaction is one that will go in two direc- 
tions depending upon the conditions of the reaction. Since 
most chemical reactions are reactions between free ions, if the 
ions are neutralized or removed, the reaction goes to com- 

17. Reactions go to completion when one of the products 
formed is (1) a precipitate, (2) a gas, or (3) water. 

18. The common-ion effect is caused by adding an ion iden- 
tical with one of the ions of a compound in solution. It results 
in a partial precipitation of that compound. 


Group A 

1 . To what questions did Arrhenius seek an answer? 

2. (a) What is an electrolyte? (b) What three classes of 
compounds are electrolytes? (c) Give three examples of non- 

3. How could you find out whether a solution contained 
an electrolyte? 

4. Explain in terms of the dissociation theory what hap- 
pens when NaCl is dissolved in H^O. 

5. Name two ways in which ions differ from atoms. 

6. In terms of the theory of dissociation define (a) an acid, 
(b) a base. 

7. (a) On what do the properties common to bases de- 
pend? (b) to acids? 

8. Why do we always use dilute acid in preparing hy- 
drogen by displacement of hydrogen from that acid by a 

9. Why will a thoroughly insoluble hydroxide not turn 
pink litmus blue? 

10. Explain neutralization in terms of ions. 

1 1 . Define strong acid and strong base in terms of the dis- 
sociation theory. 

12. What are (a) three common strong acids, (b) two com- 
mon weak acids, (c) three common weak bases, and (d) two- 
common strong bases? 


13. A solution has a pn value of six. What does this mean? 

14. Normal blood has a pn of about 7.3. Is it acid or alka- 

15. How does the degree of dissociation of NH 4 OH com- 
pare with that of NaOH in equivalent solutions? 

16. (a) What happens to a metal atom when it becomes an 
ion? (b) Explain the dissociation of NaCl in terms of elec- 

17. Explain the difference in physical and chemical prop- 
erties of the potassium atom and the potassium ion. 

18. By means of diagrams, show the difference between the 
chloride ion and the chlorine atom. 

19. Change the following equations into ionic equations. 
Consult the table of solubilities on page 243. 

a) AgNO 3 + KC1 -* AgCl + KNO 8 

b) 2NaCl + H 2 SO 4 - 2HC1 + Na 2 SO 4 

c) BaCl 2 + H 2 S0 4 - BaSO 4 + 2HC1 

20. (a) What is a concentrated acid? (b) Would a concen- 
trated acid necessarily be a strong acid? Explain. 

21. Insoluble bases are very weak bases. Explain. 

22. Complete and balance the following equations. 

a) Na 2 C0 3 + CaCl 2 -> f) Pb(NO 3 ) 2 + NaCl - 

b) Zn + H 2 S0 4 - g) BaCl 2 + K 2 SO 4 -> 

c) AgNO 8 + NaBr -> h) NaNO 3 -f CuCl 2 -> 

d) NaCl + H 2 SO 4 -> i) Pb(NO 3 ) 2 + H 2 S -> 

e) NaCl + Zn(NO 3 ) 2 -> j) NaOH + NH 4 C1 - 

23. Which of the reactions in exercise 22 are reversible and 
which go to completion? 

24. (a) Which of the following salts are neutral in solution? 
(b) Which give an acid reaction? (c) basic reaction? (d) Ex- 

(1) NaCl 1(3) A1C1 3 l/ (5)Na 3 BO 8 (7) ZnSO 4 
Na 2 CO 3 (4) KNO 3 (6) K 2 SO 4 J) Na 2 SO 8 

25. What is the percentage composition of sodium acetate, 

NaC 2 H 3 2 ? 


26. Explain the electrolysis of H 2 O in terms of ions and 

Group B 

27. Why does cold dilute H 2 SO 4 attack Zn, although cold 
concentrated H 2 SO 4 does not? 

28. Dry cobalt chloride (CoCl 2 ) is blue in color. A solution 
of this salt in water is pink. Explain why this color change 

29. In what three respects has the original theory of disso- 
ciation been modified? 

30. Concentrated H 2 SO 4 is a poorer conductor of electricity 
than dilute H 2 SO 4 . What does this indicate with reference to 

31. Explain how the study of the freezing points of solutions 
led to the theory of dissociation. 

32. (a) State the law of mass action, (b) Explain how it 
works, using a suitable reaction as an example. 

33. What is a dipole? Illustrate your answer. 

34. What weight of "Prestone," C 2 H 4 (OH) 2 , should you 
use in a 3| gallon auto radiator to protect it against 5F. 
weather? One mole of C 2 H 4 (OH) 2 lowers the freezing point 
of one liter of water from 0C. to -1.87C. Density of 
C 2 H 4 (OH) 2 is 1.13. 


1. Write a report or organize a class discussion on the im- 
portance of pn measurements to a soil (agricultural) chemist 
or to a medical research man. Consult your teacher of agri- 
culture or your family doctor. 

2. Construct a laboratory setup as shown on page 239. Plate 
some copper onto a piece of steel. The steel object should be 
the cathode and the copper strip the anode. The electrolyte 
should be a copper sulfate solution. 

3. Take an old radio tube apart and demonstrate to the 
class how electrons are liberated ifrom the heated filament, and 
how the flow of electrons passes through the grid toward the 
positive plate. Consult your teacher of physics or a radio 
engineer in connection with this project. 




The men of experiment are like 
the ant: they only collect and use; 
the reasoneis resemble spiders who 
make cobwebs out of their own sub- 
stance. But the bee takes a middle 
course; it gathers its material from 
the flowers of the garden but trans- 
forms it by a power of its ozvn. 
Francis Bacon (1561-1626) in No- 

The electron structure of ammonia. At high temperatures, nitro- 
gen combines with hydrogen, forming ammonia, NH V Nitrogen acts 
like a nonmetal with a valence of three. Since its outer ring is short 
three electrons, it combines with three hydrogen atoms, iorming a 
nonionic compound. ) 

The three-dimensional, or cubic, diagram (Fig. 54) is another 
method of showing the arrangement of electrons around the nu- 
cleus of an atom. This diagram represents the formation ot ammonia. 
Note that the electrons of the three hydrogen atoms are shared by 
the nitrogen atom in forming the ammonia molecule. Since NH, 
bears two unshared electrons, it can accept, or combine with a pro- 
ton (H+) to form NH 4 +. Hence, it is a base. 

How ammonia is prepared in the laboratory. When ammonium 
chloride is heated with calcium hydroxide, ammonia gas is liber- 
ated and may be collected by the displacement of air (Fig. 55) . 

2NH 4 C1 + Ca(OH) 2 -> CaCl 2 + 2NH 3 1 + 2H 2 O 

Since the only function of the ammonium chloride is to furnish the 
ammonium group, NH 4 , almost any other ammonium salt may be 




Ammonia NH 3 

Fig. 54. Formation of ammonia, showing electron transfer* 

substituted for it. Since any soluble base can supply the OH radical, 
we can use any soluble base instead of calcium hydroxide. The gen- 
eral method for preparing ammonia is, therefore, by the reaction of 
an ammonium salt with any soluble base. 

Physical properties of ammonia. Priestley collected ammonia over 
mercury, because it is soluble in water. At room temperature 1 vol- 
ume of water will dissolve more than 700 volumes of this gas, NH. { . 

The extreme solubility of ammonia may be shown by the ammonia 
fountain. This setup consists of a flask and a tube (Fig. 56) . The 
flask is filled with dry ammonia gas, and inverted over water. As 
soon as the tube enters the water, some ammonia dissolves in it, 
reducing the pressure inside the flask. The air pressure outside then 
forces the water up into the tube, and as it issues from the top of 
the tube it forms a fountain of ammonia water. Why? 

The characteristic pungent odor of ammonia was known to 
Priestley. It is reported that as he bent over the fireplace where he 
prepared the gas by the method we use today, its vapor made his 
eyes fill with tears and drove the occupants of his house out of doors. 

Ammonia gas is about half as heavy as air. It may be readily lique- 
fied, using only 70 pounds of pressure per square inch at ordinary 
temperatures. The colorless liquid NH 3 is kept in steel cylinders and 
shipped in tank cars. 

Fig. 55. Laboratory preparation of ammonia. 
Why is the generating tube tilted downward? 

Fig. 56. The ammonia fountain. 





Chemical properties of ammonia. Ammonia reacts with water, 
forming ammonium and hydroxyl ions. The latter ions account for 
the basic characteristics shown by the water solution of ammonia, 
aqua ammonia. 

NH 3 + H 2 * (NH 4 )+ + (OH) - 

Ammonia gas does not burn in air, but it does burn in pure oxy- 
gen with a pale greenish flame. 

4NH 3 + 5O 2 - 6H 2 O + 4NO 

Dry ammonia does not unite with dry hydrogen chloride, but the 
presence of a trace of water causes the two to combine, forming a 
white cloud that settles out as ammonium chloride powder. 

NH 3 + HC1 <=t NH 4 C1 

Ammonia unites with the very active metals. When passed over 
hot magnesium, for example, magnesium nitride and hydrogen are 

2NH 3 + 3Mg -> 3H 2 + Mg 3 N 2 

Ammonia is a very useful compound. Liquefied ammonia gas, or 
anhydrous ammonia as it is called, is a widely used compound. In a 
recent normal year, more than 2,000,000 tons were produced in this 
country for use in the chemical industries. 

More than 75 percent of all chemical nitrogen products are pro- 
duced from ammonia. Ammonia is the chief raw material from 
which nitric acid and nitrates, or salts of nitric acid, are made. Com- 
bined with acids to form ammonium sulfate, (NH 4 ) 2 SO 4 , ammo- 
nium nitrate, NH 4 NO 3 , and monoammonium phosphate, NH 4 H 3 PO 4 , 
ammonia is one of the chief sources of nitrogen in fertilizers. Urea, 
used as a fertilizer, is also made from ammonia. The manufacture 



In recent years, there has been an increasing use of anhydrous 
ammonia as a primary fertilizer. Sometimes the ammonia is added 
to irrigation water which effectively spreads it throughout the soil 
A more widespread practice is to inject liquid anhydrous ammonia 
directly into the soil through tubes mounted on a plow-like ap- 
plicator. Upon release, the ammonia, which has been kept under 
pressure, reverts to a gas and unites chemically with the soil par- 
ticles. Greatly increased crop yields have resulted from treating 
the soil with anhydrous ammonia. 

Ammonia is widely used as a refrigerant. Dry ammonia gas is used 
extensively in commercial refrigeration because of the ease with 
which it can be liquefied. In making artificial ice, ammonia gas is 
placed m a closed system of pipes and coils. By means of a pump, 
the gas is compressed until it changes to a liquid; the heat evolved 
in the process is removed by a spray of cold water. Then, as cold 
hquid ammonia, it is passed through pipes into the freezing cham- 
ber, which contains brine or a water solution of calcium chlo- 
ride, C_jaCvd 9 . 

As fast as the liquid ammonia enters the pipes in the chamber 
through a needle valve, it expands suddenly and vaporizes as a result 
ot the reduced pressure, and in so doing absorbs a great amount of 
heat from the brine. So cold does the brine become that the pure 
water m the tanks that are immersed in the brine changes to blocks 
of ice. The ammonia gas returns to the pumping chamber, where it 
is again compressed for reuse. The process is continuous, and the 
same ammonia gas is used over and over again. 

Ammonia gas also helps purify water. One method of water puri- 
fication of increasingly wide use employs both chlorine and ammo- 
nia. This method is known as the chloramine process and depends 

cold water 

liquid ammonia-*- 

1 - . =-:" " .-.-.-^" * 


Fig. 57. Ammonia refrigeration equipment. 

on the fact that when ammonia and chlorine react, chloramine, 
NH 2 C1, is formed. One of the reactions is: 

2NH, + C1 2 -> NH 2 C1 -f NH 4 C1 

Chloramine is a very effective killer of bacteria, or bactericide, but 
is less active chemically than chlorine and produces less of the typical 
taste of chlorinated water. The process is somewhat less expensive 
than chlorination, and is especially well adapted to communities 
whose water supplies have musty or swampy tastes. 

Preparation and properties of ammonium hydroxide. Care must 
be taken to distinguish between ammonia, NH 3 , and the ammo- 
nium ion, NH 4 +. In the laboratory preparation of ammonia, the ni- 
trogen compound first formed might be thought of as ammonium 
hydroxide, which breaks down into ammonia and water. 

2NH 4 C1 + Ca(OH) 2 -> CaCl 2 -f [2(NH 4 )+ + 2 (OH)-] 

2NH 3 + 2H 2 

A water solution of NH 3 is often called ammonium hydroxide even 
though the compound probably does not exist. The solution really 
consists of some dissolved NH 3 , some NH 4 ions, and some OH ions. A 
saturated solution of ammonia is lighter than water (sp. gr. 0.88) , 
and contains about 36 percent NH 8 by weight at room temperatures. 
The ammonia may be expelled by boiling. Household ammonia is 
a water solution of NH 3 containing about 6 percent oleic acid. 

A water solution of ammonia is a weak base, since NH 3 reacts with 
water to form only a few (one percent) hydroxyl ions. Because it is 
a base, it dissolves grease and hence removes dirt. Since this ammo- 
nia water gives off vapor, or volatilizes, rapidly and completely with- 
out leaving a solid, it is useful as a household cleansing agent. 




The ammonium radical and ammonium salts. The ammonium 
radical, because of its positive valence, is considered metallic. Its 
presence may be detected by adding the salt to be tested to a base, 
such as sodium hydroxide, and heating the mixture. If the sus- 
pected substance is an ammonium compound, ammonia gas will be 
liberated, and can be identified by its odor. Moist pink litmus paper 
held in the gas turns blue. Why must the litmus be wet? 

(NH 4 ) 2 S0 4 4 2NaOH -+ Na 2 SO 4 4- 2H 2 O 4- 2NH 3 T 

One of the most common of the ammonium salts is ammonium 
chloride, or sal ammoniac. This salt, first produced in Egypt, was 
known to the early alchemists. It is a white, crystalline substance, 
readily soluble in water. It is decomposed by heat into two gases, 
NH 3 and HC1, which reunite on cooling. 

,, NH 4 C1 *=> NH 3 4- HC1 

Sal ammoniac is used extensively in dry batteries as an electrolyte 
and in soldering in which the hot iron dissociates the salt. The hydro- 
gen chloride arid ammonia liberated remove the rust that covers the 
surface of the metal to be soldered, the hydrogen chloride by dissolv- 
ing the rust, and the ammonia by reducing it. 

Fritz Haber makes the synthesis of ammonia a commercial suc- 
cess. Minor quantities of ammonia are still made by an old process 
the destructive distillation of coal. This process, in which ammonia 
is produced as a valuable byproduct, is discussed on pages 862 363. 
Synthetic methods have superseded the method of preparing am- 
monia from coal. The most successful synthesis of ammonia is based 

A compressor used in the manufacture of ammonia. The hydogen is made from 
natural gas produced in nearby oilfields. 

Shell Chemical Corporation 


on the process first worked out on a commercial basis by Fritz Haber 
(ha'ber) in 1913 and known as the Haber process. 

The first real test of this great achievement came during World 
War I, partly as a result of the desperate need of the German gov- 
ernment for nitrogen compounds. Haber's process made agriculture 
in blockaded Germany independent of Chile saltpeter and also gave 
the German military machine a new source of nitrates for high ex- 
plosives. The Haber process enabled Germany to fight hunger, and 
stave off defeat much longer than the Allies expected. Fritz Haber 
was later forced into exile by the Nazis and died in Switzerland. 

The Haber process. This process, with its many modifications, is 
the most important single process for the fixation of nitrogen, that is, 
the combining of the free nitrogen of the air with other elements to 
form useful compounds. The process is based upon what appears to 
be a very simple reaction, the union of hydrogen and nitrogen gases. 

The nitrogen and hydrogen for this reaction are obtained from 
air, coke-oven gas, water gas, natural gas, and some petroleum refin- 
ery gases. Although this chemical reaction has been known for a 
long time, it was not industrially practicable until the reversible 
reaction could be controlled. The laws of chemical equilibrium had 
to be used, so that some of the ammonia gas formed would not be 
immediately decomposed into its elements. 

What factors can be used to control the point of equilibrium? 
Most chemical reactions are reversible. In the synthesis of ammonia, 
the first equilibrium ratio was about two percent ammonia to 98 per- 
cent of a mixture of nitrogen and hydrogen. In other words, most of 
the ammonia formed during the union of nitrogen and hydrogen de- 
composed into its constituent gases. 

The point of equilibrium, however, can be controlled to some 
extent. The factors that help to control it are (1) temperature, 
(2) pressure, and (3) concentration of the substances involved in 
the reaction. Catalytic agents increase the speed of a reaction and, 
hence, enable the point of equilibrium to be reached more rapidly. 
It is not probable that catalysts alter the point of equilibrium. 

Thus, in the case of the preparation of ammonia in America by a 
modified Haber process, the reaction 

3H 2 + N 2 -* 2NH 3 T + 24,000 calories 

has been forced to go to the right, producing as much as 30 percent 
NH 3 , instead of only two percent of the theoretical yield. 



\itrogen Division, Allied Chemicul rind Dye Corporation 

Aerial view of a large plant for the synthesis of ammonia and the manufacture 
of fertilizer. 

The conditions that made the American process a success were 
(1) the use of a specially prepared iron oxide as catalyst, (2) a tem- 
perature of about 475C., (3) a reaction pressure of about 300 atmos- 
pheres (atrn.) , that is, 300 times atmospheric pressure, and (4) a 
rapid removal of the NH 3 formed. Since heat is evolved during the 
synthesis of ammonia, the higher the temperature, the less the yield, 
and consequently too high a temperature is avoided. Since there is 
a diminution of volume (three volumes of hydrogen unite with one 
volume of nitrogen, forming only two volumes of ammonia) , the 
yield is increased by an increase in pressure. The manufacturing 
conditions represent the most effective compromise between largest 
yield, shortest time, and most profitable rate. 

Another synthetic process for making ammonia. In 1916 the 
United States built a large nitrogen-fixation plant at Muscle Shoals, 
Alabama, where ammonia was to be prepared by the cyanamide proc- 
ess. This project was later abandoned, although the cyanamide 
process at Niagara Falls, Canada, has been very successful as a source 
of calcium cyanamide. The chief chemical changes that take place 
in the cyanamide process are: 

1) The formation of calcium carbide by heating lime, CaO, and 
coke, C, in an electric furnace. 

CaO + 3C - CO -j- CaC 2 (calcium carbide) 


2) The union of calcium carbide with free nitrogen, forming 
calcium cyanamide. 

CaC 2 + N 2 > C + CaCN 2 (calcium cyanamide) 

3) The addition of steam to cyanamide, forming ammonia. 

CaCN 2 + 3H 2 O - CaCO 8 + 2NH 3 f 

In dry form, crude, powdered calcium cyanamide containing about 
60 percent CaCN 2 is sold as a fertilizer under the name of "Cyana- 
mid." All of its nitrogen is available as a plant food. 


Berl, E. "Fritz Haber." Journal of Chemical Education, May, 
1937, pp. 203-207. A short biography. 

Jaffe, Bernard. Men of Science in America, pp. 307-330. 
Simon & Schuster, New York, 1944. The development of our 
ideas regarding reversible reactions is tied up with the contri- 
bution of America's greatest theoretical scientist, J. Willard 
Gibbs. His life and work are described here. 

Ross, William H.; Adams, J. Richard; Yee, J. Y.; and Whit- 
taker, Colin W. "Preparation of NH 4 NO 3 for Fertilizer Use." 
Industrial and Engineering Chemistry, Dec., 1944, pp. 1088- 

Slosson, Edwin E., Creative Chemistry, pp. 14-36. D. Apple- 
ton-Century Co., New York, 1920. "Nitrogen, Preserver and 
Destroyer ol Lite." 


1. When a substance which ionizes goes into solution, the 
change that takes place is reversible. The substance dissociates 
into ions; these ions unite, re-forming the original substance. 
Even when equilibrium is established, this reversible reaction 
continues. Change is no longer apparent, however, because 
the rate of dissociation is the same as the rate at which the ions 
in the solution reunite, forming the original substance. The 
solution is in a state of dynamic equilibrium. 

2. The point of equilibrium ot a reversible reaction can be 
controlled to a certain extent by (1) temperature, (2) pres- 
sure, and (3) concentration of the substances involved. Cata- 
lytic agents increase the speed of a reaction and, hence, enable 
the point of equilibrium to be reached more rapidly. 



Group A 

1. Describe the laboratory preparation of NH 3 . Use a 
labeled diagram. 

2. (a) How did Priestley first collect ammonia? (b) Why 
did he use this method? 

3. (a) What is the general method of preparing NH 3 ? 
(b) Write two equations illustrating two ways of preparing 
NH 3 from NH 4 C1. 

4. Show, by a simple experiment, the extreme solubility of 
NH, in H 2 O. 

5. Devise a simple experiment to determine whether NH 8 
is lighter or heavier than air. 

6. Does NH 3 burn? Explain. 

7. (a) By means of an equation, give the chief chemical 
property of NH 3 . (b) What type of reaction is this? (c) How 
can you make the reaction go to the right? (d) to the left? 

8. Complete and balance the following equations: 

a) NH 3 + HC1 - c) (NH 4 ),S0 4 + NaOH-> 

b) NH 3 + Mg -> d) NH 3 + H 2 SO 4 


9. What are five uses of NH 3 ? 

10. What substances are present in a water solution of NH 3 ? 

11. (a) Is a solution of NH 3 in water a strong base? (b) Ex- 
plain your answer. 

12. (a) Write an ionic equation showing the neutralization 
of HC1 by NH 4 OH, (b) also the neutralization of H 2 SO 4 
and (c) of HNO 3 by the same base. 

13. Does dry NH 3 affect litmus? Explain. 

14. (a) In what group of elements is the ammonium radical 
placed? (b) Explain. 

15. What are two commercial methods of preparing NH 3 ? 

16. What was the difficulty that confronted manufacturers 
who attempted to make NH 3 by direct synthesis of its ele- 

17. What four conditions are met in the improved Haber 


18. (a) Describe the cyanamide process for making NH 8 . 
(b) What are the three equations? 

19. Complete and balance the following equations: 

a) NH 3 + HNO 3 -> b) NH 3 + H 3 PO 4 -> 

20. "Spirits of hartshorn" was the name applied to am- 
monia water prepared by the alchemists by heating the horns 
of deer. What elements must have been present in deer horns? 

21. In what two ways do liquid ammonia and aqua am- 
monia differ? 

22. What simple test will distinguish NH^Cl from Nad? 
Both of these compounds are white, soluble salts. 

23. Compare the laboratory methods of collecting NH 3 , N 2 , 
HC1, 2 , C1 2 , and H 2 . 

24. Explain the action of liquid NH 3 on dry litmus paper. 

25. Why is ammonia water called the volatile alkali? 

26. Bottles of household ammonia were formerly closed 
with rubber rather than with glass or cork. Why? 

27. Why is ammonia water, rather than lye, used in re- 
moving grease spots from clothing? 

28. Compare the ease with which NH 3 , N 2 , C1 2 , and H 2 are 

29. Describe the preparation of artificial ice by means of 
NH 3 . (b) What property of NH 8 makes it useful in this proc- 
ess? (c) Of what use is the CaCl or brine solution? 


30. What weight of NH 4 C1 is necessary to make 340 tons 
of dry liquid NH 3 ? 

31. Determine the percentage of nitrogen in monoammon- 
ium phosphate. 

32. (a) What is chloramine? (b) How is it prepared? 

Group B 

33. Bottles of ammonia water and hydrochloric acid are 
placed within a few inches of each other and their stoppers re- 
moved. White fumes are seen. Explain. 


34. Explain how you would determine the strength of a bot- 
tle of household ammonia by titration. Give details. 

35. How can a knowledge of the laws of equilibrium be 
used in making the preparation of NH 8 by the Haber process 
more efficient and more economical? 

36. When water is added to magnesium nitride (Mg 3 N 2 ) , 
the odor of NH 3 is detected. Write an equation to explain this. 

37. CaCl 2 unites with NH 3 , forming CaCl 2 8NH 3 . Explain 
whether you could use CaCl 2 in drying ammonia gas. 

38. Account for the odor of NH 3 around heaps of garbage 
and manure. 


1. Take a dry cell apart, and show the various substances 
that are used in its manufacture. Prove the presence of NH 4 C1 
by a chemical test. What is the purpose of the NH 4 C1? 

2. Write a two- or three-page report on the freedom of the 
man of science in a democracy compared with the enslavement 
of science in totalitarian states. Use Haber and, perhaps, Lang- 
muir as examples. 

3. Set aside 2 ml. of household ammonia in an open test 
tube for two weeks. Test for the presence of ammonium hy- 
droxide with pink litmus paper at the end of the time. Explain 
the results. 




. . . For nitrogen plays a double role 
in human economy. It appears like 
Brahma in two aspects, Vishnu the 
Preserver and Siva the Destroyer. 
E. E. Slosson, 1919 

The revolution brought about by man-made nitrogen com- 
pounds. Haber's successful synthesis of ammonia widened man's 
control over nature by making him tess dependent for his raw mate- 
rials on supplies present in limited or faraway areas of the earth's 
surface. Synthetic ammonia was soon converted into nitric acid. 
This, in turn, gave mankind an unlimited supply, not only of ferti- 
lizers, but also of high explosives. The synthesis of nitrogen com- 
pounds from air and water was a bloodless revolution whose conse- 
quences touched the lives of half the people of the world. 

Preparation and properties of nitric oxide. Not only does nitrogen 
unite with hydrogen at high temperatures, but it combines with oxy- 
gen also when an electric spark is passed through a mixture of the 
two gases. Cavendish made this discovery in 1770 when he passed an 
electric spark through mixtures of hydrogen and air in his synthe- 
sis of water. Soon afterward Priestley made a thorough study of the 
compound formed by the direct union of nitrogen and oxygen. 

N 2 + O 2 2NO (nitric oxide) 

Priestley also prepared nitric oxide by the action of copper on 
dilute nitric acid, collecting the gas by the displacement of water, 




as shown in the illustration below. This is the laboratory method 
used today. 

3tu + 8HNO 3 -> 3Cu(NO 3 ) 2 4- 4H 2 O + 2NO T 

Nitric oxide is a poisonous, colorless gas, very slightly soluble in 
water, and about as heavy as air. Chemically, it is very active. When 
exposed to air or oxygen, it oxidizes at once to nitrogen dioxide, NO 2 . 
This reaction produces much heat; that is, it is decidedly exothermic. 

Nitrogen dioxide. The equation representing the oxidation of 
nitric oxide to nitrogen dioxide is: 

2NO + O 2 

2NO 2 

Nitrogen dioxide is reddish brown in color, heavier than air, very 
soluble in water, and easily liquefied. Its fumes are irritating and 

In preparing nitric oxide by means of copper and nitric acid, the 
gas first seen in the generating bottle is brown nitrogen dioxide 
rather than the colorless nitric oxide. Actually NO is formed first, 
but combines with the oxygen in the air of the generating bottle to 
form the brown NCX. When the mixture of the NO and NO 2 passes 
through the delivery tube into the collecting bottle, the NO 2 dis- 
solves in the water. Only the colorless nitric oxide displaces the 
water in the bottle. 

As its temperature is lowered, nitrogen dioxide gradually changes 
into nitrogen tetroxide gas, N 2 O 4 . The reaction 2NO 2 ^ N 2 O 4 is a 
reversible one. Above 140C., the reaction goes to completion, as 
shown by the arrow pointing to the left. At room temperature, the 
gas obtained is a mixture of the reddish brown NO 2 and the color- 
less N 2 O 4 . At very low temperatures, NO 2 changes completely into 
N 2 4 / 

The arc process of making nitric acid from air and water. The 
union of nitrogen with oxygen in an electric arc, or by the action of 
an electric spark, was used as the basis of many commercial ventures 
attempting to produce synthetic nitric acid. In the beginning, they 

nitric oxide 


Fig. 58. Laboratory prepara- 
tion of nitric oxide. What other 
laboratory setup does this re- 

copper wire 
inside glass 


Fig. 59. Laboratory preparation of nitrogen dioxide by the arc process. How can 
you tell whether NO 2 is formed? 

all failed because when a mixture of nitrogen and oxygen was thus 
treated, the yield of nitric oxide was very small. Because the reac- 
tion is reversible, practically all the nitric oxide formed at first was 
decomposed into its original elements. 

N 2 4- O 2 ?= 2NO T 

A careful application of the laws of chemical equilibrium finally re- 
sulted in larger yields of nitric oxide. This gas combined with the 
oxygen of the air to form nitrogen dioxide which was then dissolved 
in water to form nitric acid, HNO,. 

2NO + 2 

3N0 2 + H 2 

2NO 2 

> 2HN0 3 4- NO 

The Ostwald process, another commercial method of making 
nitric acid. The arc process because of its higher cost has been super- 
seded by the Ostwald process. No sooner had the synthesis of ammo- 
nia been successfully carried out than Wilhelm Ostwald (ost'valt) , 
a chemist who helped Arrhenius establish the theory of ionization, 
showed how ammonia gas could be converted into nitric acid on an 
industrial scale. 

The Ostwald process consists of oxidizing ammonia gas in the pres- 
ence of a catalyst, which consists of a specially-prepared platinum or 
platinum-rhodium gauze heated to a red heat. The two oxides of 
nitrogen are formed during the process, but the final products may 
be represented by the following equation: 

NH 3 + 2O 2 -> HN0 3 + H 2 O 

Using a pressure of only about six atmospheres, about 95 percent 
of the ammonia is converted into a water solution of nitric acid of 
about 50 percent concentration. 

Nitrogen fixation. The change of free nitrogen into ammonia by 
the Haber process and the change of free nitrogen from the air into 
nitric acid by the arc process are examples of nitrogen fixation. 


gauze" 1 combustion 



H 2 O 

Fig. 60. Laboratory preparation of NOj by the Ostwald process. 

Nitrogen fixation is the changing of free nitrogen from the air into 
useful compounds. During electric storms, some nitric acid is formed 
in the atmosphere by a natural arc process. 

Nitrogen-fixing bacteria help the farmer. Most farm crops use 
part of the valuable salts in soil. On the other hand, certain crops 
such as peas, beans, and alfalfa, actually enrich the soil in which 
they are grown. Chemistry explains the fertility of such soil in the 
following way: Plants require nitrogen in the form of nitrates. 
These nitrates are soluble in water and can, therefore, be absorbed 
by the root hairs of plants by osmosis, a process by which liquids and 
gases pass through semipermeable tissues. On the roots of plants such 
as peas, beans, and alfalfa are little nodules, inside which live bil; 
lions of nitrogen-fixing bacteria (Rhizobium) . These bacteria and 
several others have the ability to change the free nitrogen, present 
in porous soil, into soluble nitrogen compounds, particularly ni- 
trates, that plants use in building living tissues. When such crops 
are plowed under, this "green manure" adds nitrogen to the soil. 

The nitrogen cycle nitrogen compounds break down into free 
nitrogen. Various other bacteria break down nitrogen compounds in 
the soil into simpler compounds and even return considerable quan- 
tities of free nitrogen to the air. Such bacteria, called denitrifying 
(de-ni'trMI-ing) bacteria, cause the loss of nitrogen from the soil and 
thus complete the extremely important nitrogen cycle. These bacteria 
are responsible, in part, for the rapid decay of nitrogen-containing 

These pea roots are covered 
with nodules containing nitro- 
gen-fixing bacteria. 


organic wastes. They are used widely in the treatment of sewage (see 
pages 222-223) . This series of changes is referred to as the nitrogen 

How nitric acid is prepared in the laboratory. The laboratory 
preparation of nitric acid follows the general method for preparing 
an acid. Sodium nitrate, mixed with concentrated sulfuric acid, is 
heated gently in a glass retort. Nitric acid is formed, which, having 
a lower boiling point than sulfuric acid, evaporates. It is then con- 
densed into a colorless liquid by cooling, as shown in the illustration 
below. The equation for the preparation of nitric acid is: 

NaNO 3 + H 2 SO 4 -> HNO 3 + NaHSO 4 (sodium hydrogen sulfate) 

This method is also one of the commercial processes used today in 
making about 10 percent of the nitric acid consumed by the world's 
industries. Synthetic NaNO 3 supplies the nitrate. 

Nitric acid was known to the alchemists more than 1000 years ago. 
Geber (ga'ber) , an Arabian physician and alchemist, prepared it 
about A.D. 800. It was called aqua fortis, meaning strong water. 

Physical properties of nitric acid. When pure, nitric acid is a color- 
less liquid. Its water solution, containing 68 percent nitric acid by 
weight (that is, the concentrated nitric acid of commerce) , has a 
specific gravity of 1.4, and boils at 120C. The concentrated acid 
fumes strongly. The yellowish appearance of the nitric acid prepared 
in the laboratory is caused by the presence of nitrogen dioxide, 
formed by the partial decomposition of the nitric acid during the 

Chemical properties of nitric acid. Because nitric acid mixes with 
water in all proportions and dissociates almost completely, thus pro- 
ducing large quantities of hydrogen ions, it is a strong acid. Nitric 
acid is unstable. In sunlight or when heated, it decomposes into 
water, oxygen, and nitrogen dioxide. 

4HN0 3 -- 2H 2 + 4NO 2 T + O 2 T 
Fig. 61. Laboratory preparation of nitric acid. By what process is the acid collected? 

cold water 

nitric acid 


A glowing splint inserted in the vapors of boiling concentrated 
nitric acid catches fire, thus showing that oxygen is present in the 
vapors. When an element, such as nitrogen in HNO 3 , is in a very 
highly oxidized state, that is, it has a high positive charge, the com- 
pound is a strong oxidizing agent. 

The action of nitric acid on metals illustrates its oxidizing pow- 
ers. When hydrochloric acid reacts with many of the metals, hydro- 
gen is liberated even if concentrated acid is used. On the other hand, 
dilute nitric acid acts on a metal, forming water instead of hydrogen. 
Nitric oxide gas is also formed. In fact, one of the methods used to 
prepare nitric oxide depends on this action of nitric acid. 

3Cu + 8HNO 3 -> 3Cu(NO 3 ) 2 + 4H 2 O + 2NO j 

When concentrated nitric acid reacts with a metal, nitrogen 
dioxide, instead of nitric oxide, is formed. Brown NO 2 is produced 
by the oxidation of NO to NO.. 

Cu + 4HN0 3 - Cu(N0 3 ) 2 + 2H 2 O + 2NO 2 1 

The electron theory explains the oxidizing power of nitric acid as 
follows: oxidation is a loss of electrons; therefore, a substance such 
as chlorine, which borrows electrons, is a good oxidizing agent. 
HNO 3 may be thought of as containing N+++++ , which borrows elec- 
trons from the copper in the reaction above, changing to N++ (in 
NO) . Thus the copper is said to be oxidized and the nitrogen re- 

Nitric acid attacks proteins, forming a yellow compound. The 
yellow coloration produced on the skin by nitric acid is caused by 
this chemical reaction (see also Chapter 36) . Nitric acid oxidizes 
both cotton and wool, destroying most of the strength of the fibers. 

Aqua regia, the acid mixture that dissolves gold and platinum. 
Nitric acid, when mixed with hydrochloric acid, oxidizes the latter, 
liberating atomic chlorine. This mixture of nitric and hydrochloric 
acids as in the equation below is called aqua regia, or royal water. 

3HC1 + HNO 3 - 2H 2 O + NO + 3C1 

This chlorine reacts with gold, forming soluble gold chloride. 

Au + 3C1 -> AuCl 3 

Both gold and platinum are insoluble in any one of the three com- 
mon strong acids, but they are soluble in aqua regia. 


FeSO 4 + nitrate solution -j^nar i 

jjjji&r concentrated 

brown ring--"^^ H 2 SO 4 

Fig. 62. The brown-ring test for a nitrate. 

How do we test for the nitrate ion? The nitrates, salts of nitric 
acid, are all soluble in water, are decomposed by heat, and may 
be detected by the brown-ring test. This test is made by adding a 
small amount of freshly prepared ferrous sulfate solution to the solu- 
tion suspected of containing a nitrate. Concentrated sulfuric acid is 
then carefully poured down the side of the test tube in such a way 
that it collects at the bottom without mixing with the solution. If 
a brown layer forms between the heavy sulfuric acid and the solu- 
tion floating on top, a nitrate is present. The sulfuric acid pro- 
duces nitric acid by reacting with the nitrate, and the nitric acid 
in turn reacts with the ferrous sulfate, forming a brown compound, 
FeSO 4 NO. 

Gunpowder, the earliest explosive. The inactive element nitrogen 
does not unite easily with other elements. And after it does, the 
unions so produced are very unstable. In fact, such unions are so 
extremely unstable that on the slightest provocation * the nitrogen 
breaks away with a bang! Most explosives, except those based on 
nuclear fission, depend upon this fact. 

Nearly all nonfission explosives contain either nitrate (NO 3 ) , or 
nitro (NO 2 ) , radicals. In addition, some explosives contain ammo- 
nium radicals, NH 4 . When compounds that contain nitrate or nitro 
radicals are mixed with other compounds that can easily use the oxy- 
gen of these unstable radicals, an explosive is the result. In certain 
cases, the compound containing the nitrate or nitro radicals actu- 
ally supplies the means of its own destruction by furnishing the ele- 
ments that can use the oxygen readily. When something, such as a 
shock, starts the reaction, the unstable nitrate or nitro radicals re- 
lease their oxygen for combination with other elements and liberate 
free nitrogen gas. Nearly always, most of the other products of the 
reaction are gases also, and because of the high temperatures pro- 
duced, terrific pressures result. 

The earliest explosive was made thousands of years ago by 
the Chinese. How the black powder came to be produced is not 
known, but it was made then about as it is now, by mixing approxi- 
mately 15 parts by weight of potassium nitrate, KNO 3 , with three 
parts of sulfur, and two parts of powdered charcoal. The resulting 


black powder, or gunpowder, explodes with terrific violence, much 
lire, and great quantities of acrid smoke. The reaction produced is: 

2KNO 3 + 3C + S - K 2 S + N s t + 3CO 2 T 

The Chinese used gunpowder in producing firecrackers and other 
kinds of fireworks for use in ceremonies and celebrations. But some- 
time within the past few centuries, no one knows for sure just when, 
gunpowder was turned to military uses, and, like so many other sci- 
entific discoveries, soon created a revolution. The foot soldier with 
his primitive musket quickly replaced the heavily-armored and be- 
pl umed knight with his lance and shield. Turrets and thick stone 
walls were no defense against powerful cannons, and the picturesque 
castle of the Middle Ages became obsolete. 

Nitrocellulose, nitroglycerin, and some detonators. Today, gun- 
powder is considered a relatively "tame" explosive. Since its inven- 
tion, chemists have produced many kinds of explosives chiefly by 
nitrating, or adding NCX or NO, radicals to, such substances as cot- 
ton, glycerin, sugar, starch, and other organic compounds. Nitrocel- 
lulose, or guncotton, was produced by Schonbein (shun'bln) in 1846 
by nitrating cotton with a mixture of nitric and su If uric acids. Nitro- 
glycerin, made by nitrating glycerin, a common byproduct of the 
manufacture of soap, was produced in 1847. Both of these compounds 
are more powerful explosives than gunpowder, but both are much 
more sensitive to shock and, hence, explode much more easily. 

In 1888, Alfred Nobel, a Swedish inventor, produced dynamite 
by absorbing nitroglycerin in a fine clay, or diatomaceous earth (see 
page 500) . Dynamite is much less sensitive to shock and, hence, 

Standard Oil Company (\'.J.) 

A workman inserts a dynam 
cartridge into a drilled h 
during the construction of 
underground pipeline. 



can be used with greater safety than nitroglycerin. American dy- 
namite usually consists of nitroglycerin absorbed in a wood meal 
that resembles fine sawdust. The substance is packed in sticks com- 
posed of parchment paper. So overcome was Nobel by the possible 
uses of his invention that he dedicated part of its profits to the es- 
tablishment of the Nobel Peace prize for outstanding contributions 
to the peace of the nations of the world, and of the Nobel prizes for 
outstanding contributions to research in physics, chemistry, medi- 
cine and physiology, and literature. Peacetime uses of explosives in- 
clude mining, building dams, and other construction work. 

Explosives are set off, or detonated, by means of a shock produced 
by even more unstable, and thus more sensitive, compounds of ni- 
trogen called detonators. Fulminate of mercury, Hg (CNO) 2 , a 
widely used detonator, is 1 1 times as sensitive to shock as trinitro- 
toluene, or TNT, and twice as sensitive as nitrocellulose. Lead azide, 
Pb (N 3 ) 2 , another detonator, is half as sensitive to shock as fulminate 
of mercury. Such substances are used in making the caps and other 
devices with which explosives are set off. 

Nitric acid has many peacetime uses. Aside from the production 
of explosives, a major industry even in peacetime, the chief use of 
nitric acid is in the production of nitrates of organic compounds, 
such as nitrocellulose and nitrobenzene. Nitrocellulose is used in 
making some photographic films, and many quick-drying lacquers 
and enamels, especially for the automobile industry. Nitrocellulose 
is used also in the production of many kinds of artificial leathers. 
Nitrobenzene is the basic raw material of the aniline, or coal-tar 
dye, industry. 

Nitric acid also furnishes the oxides of nitrogen required in the 
chamber process for the manufacture of sulfuric acid (see page 311). 
Nitrates for fertilizers and metallic nitrates are made from nitric 
acid. Sodium nitrate and ammonium nitrate are the chief fertilizers 
produced. Strontium nitrate is one of the chemicals used in pyro- 
technics, the production of fireworks, which consist mainly of flares 


Use by approximate percentage 


and shells that give off flames and smokes of various colors. A one 
percent silver nitrate solution is put in the eyes of newborn babies 
to prevent infection that may lead to blindness. 

Nitric acid plays an important role in the pickling of steel, in the 
etching of engravers' plates for printing, and in the manufacture of 
the arsenate insecticides (chiefly lead and calcium arsenates) so 
widely used against the boll weevil, and in the spraying of fruit 
trees (see page 453) . 

The nitrogen situation today represents a chemical revolution. 
It should be apparent already that the production of nitrogen com- 
pounds is a basic industry. The world's normal consumption of man- 
ufactured nitrogen compounds is many millions of tons annually. 
At one time, sodium nitrate from Chile and nitrogen compounds ob- 
tained as byproducts from the coal industry and slaughterhouses 
were the only sources of nitrogen compounds. Today, the total an- 
nual amount of fixed nitrogen produced by chemical methods is 
many times as great as the total annual consumption of both Chilean 
nitrate and all byproduct nitrogen taken together. The story of the 
nitrogen industry bears testimony to the widespread development 
of synthetic chemistry. It has changed the economic life of millions. 

There are six oxides of nitrogen. In addition to the two oxides 
of nitrogen already discussed, four other oxides are known. The com- 
plete list is: nitrous oxide, N 2 O; nitric oxide, NO; nitrogen diox- 
ide, NO 2 ; nitrogen trioxide, N 2 O^ nitrogen tetroxide, N.X) 4 , a pow- 
erful oxidizing agent; and nitrogen pentoxide, N 2 O r ,. They illustrate 
the law ot multiple proportions and the fact that nitrogen has sev- 
eral different valences. Why? 

When nitrogen trioxide, N 2 O y , is added to water, nitrous acid, a 
very unstable acid, is formed. 

N 2 O 3 + H 2 O -> 2HNO 2 

Similarly, the addition of water to nitrogen pentoxide, N 2 O 5 , pro- 
duces nitric acid. 

N 2 5 + H 2 -> 2HN0 3 

These two gases may therefore be said to be the anhydrides (with- 
out water) of nitrous and nitric acids, respectively. An acid anhy- 
dride is an oxide whose water solution is an acid. 

Nitrous oxide, or laughing gas. Priestley was the first to produce 
nitrous oxide, N 2 O, a colorless, heavy gas, slightly sweetish in odor 
and somewhat soluble in water. This was about two years before he 

A modern hospital operating 
room. Suspended from the 
table in the left foreground 
are tanks of pure oxygen, 
carbon dioxide, nitrous oxide 
and other gases ready for im- 
mediate administration when 

Ohio Chemical and Surgical hqu, 

discovered oxygen. He learned that it supported the burning of a 
candle better than did ordinary air. It decomposes rather easily into 
oxygen and nitrogen. Just before the close of the eighteenth century, 
Humphry Davy achieved fame overnight by his discovery of the 
physiological effects of this gas. He breathed four gallons of it and 
noticed its power to produce a peculiar intoxication, which included 
laughing. The poet Samuel Coleridge, as well as other distinguished 
persons, came to Davy's London laboratory to experience the thrill 
of inhaling this gas. Nitrous oxide is still prepared as Davy made it, 
by heating ammonium nitrate. 

NH 4 N0 3 -> 2H 2 + N 2 O 

In 1842 ether was used as the first anesthetic in surgery by Dr. 
Crawford W. Long, a country doctor of Georgia. William Morton's 
use of ether at the Massachusetts General Hospital in 1846 intro- 
duced this anesthetic to the medical world. Two years earlier, Dr. 
Horace Wells, a dentist of Hartford, Connecticut, had one of his 
teeth extracted after he had anesthetized himself with nitrous oxide. 

Today nitrous oxide is still used as an anesthetic in many opera- 
tions, especially those of dentistry. It is usually mixed with about 
25 percent oxygen and, in cases of more serious operations, with 
ether. This mixture of nitrous oxide and oxygen can be breathed 
for a considerable period without harmful effects on the circulatory 



system or on vital organs. Small amounts of nitrous oxide are used 
in preserving perishable foods and liquids. Easily liquefied, it is sold 
in cylinders. It is used to eject whipped cream at soda fountains. 


Conant, James Bryant. The Overthrow of the Phlogiston 
Theory. Case 2 of the Harvard Case Histories in Experimental 
Science. Harvard University Press, Cambridge, Mass., 1950. 
Gives an excellent account of Priestley's confusion between 
oxygen and nitrous oxide. 

Haynes, William. This Chemical Age, pp. 78-94. Alfred A. 
Knopf, New York, 1942. A discussion on explosives and their 
relation to the dye industry, entitled "Mars: Chemical Dic- 

Slosson, Edwin E. Creative Chemistry, pp. 37-59. D. Apple- 
ton-Century Co., New York, 1920. A very readable account of 
nitrogen and nitrogen compounds in relation to plants. 


Group A 

1. Explain the arc process of making HNO 3 . 

2. Write equations showing two methods of preparing NO. 

3. State one chemical and three physical properties of NO. 

4. Make a diagram showing the laboratory preparation 
of NO. 

5. (a) What happens when NO comes in contact with 
air? (b) Explain. 

6. Under what conditions does NO change into N.,O 4 ? 


7. Write the equation that is the basis of the Ostwald proc- 
ess for the synthesis of HNO 3 . 

8. (a) What natural phenomenon results in the formation 
of certain oxides of nitrogen? (b) Explain. 

9. (a) What is nitrogen fixation? (b) Illustrate. 

10. Make a diagram showing the laboratory preparation of 
HN0 8 . 

11. When a mixture of NaNO 3 and H 2 SO 4 is heated in a 
retort, HNO 3 is formed a little at a time, (a) What four 


substances are in the retort? (b) Which is removed by heat? 
Why? (c) Why do the other substances remain? 

12. (a) What property of HNO 3 makes it possible to pre- 
pare the acid by the laboratory method? (b) Why could not 
HC1 be used instead of H 2 SO 4 ? 

13. Write the equation for the laboratory preparation of 
HNO 3 . 

14. State four chemical properties of HNO 3 . 

15. (a) How is aqua regia prepared? (b) Its power to dis- 
solve gold results from what property? (c) What property of 
HNO 3 is shown? 

16. (a) Why does HNO 3 appear to be yellow when prepared 
in the laboratory? (b) What property of HNO rf does this color 

17. (a) What oxide of nitrogen is always formed when 
HNO 3 decomposes? (b) What oxide of nitrogen is always 
formed when HNO 3 acts as an oxidizing agent? 

18. How would you test for the presence of the nitrate ion? 

19. What are the three principal uses of HNO/ 

20. (a) Most explosives are based upon what fact? (b) What 
is a detonator? 

21. (a) What is gunpowder? (b) dynamite? 

22. Copy and complete the following: The six oxides of 
nitrogen illustrate the law of .... The anhydride of HNO 3 
is . . . , and the anhydride of HNO 2 is .... Another name for 
N 2 O is .... N 2 O was first used as ... by Dr. Horace Wells. 
N 2 O is prepared by heating .... 

23. 120 g. of NO are obtained by the action of Cu on 
HNO 3 . How many grams of Cu (NO 3 ) 2 are formed? 

24. How do you explain the fact that the reaction of 
N 2 -(- O 2 2NO, although known for more than a century, 
could not be used in the preparation of HNO. until rather 

I ... 

25. Why is N 2 O a better supporter of combustion than NO? 

26. (a) How does nature restore some of the nitrogen com- 
pounds taken from the soil by growing crops? (b) What is 
"green manure"? 


27. In 1898, Sir William Crookes, one of England's most 
eminent chemists, predicted a future world famine caused by 
exhaustion of nitrogen compound fertilizers. Why has his pre- 
diction not materialized? 

28. Make a table showing the properties of the six oxides of 

29. How would you separate Au from Cu in a copper-gold 

Group B 

30. Why cannot HNO 3 be used in preparing H 2 ? Answer 
this question in the light of the electron theory. 

31. (a) Using the electron theory, explain how HNO 3 oxi- 
dizes Cu. By inspection of the equation, state how many atoms 
of Cu are oxidized, (b) How many atoms of nitrogen are re- 
duced? (c) How does the total number of electrons lost com- 
pare with the total number gained? 

32. How could you determine experimentally whether a 
gas contained a high percentage of N 2 O 4 and a small amount 
of NO 2 , or was composed almost entirely of NO,? 

33. There is less oxygen in nitrous oxide than in nitric 
oxide, (a) Which would support burning better? (b) Explain. 

34. Nitrates are unstable in the presence of heat. What 
product would you expect Cu (NO 3 ) 2 to yield when heated? 


1. At Muscle Shoals, Alabama, there is a huge plant for 
generating electric power. This plant is capable of producing 
thousands of tons of nitrogen compounds a year by fixation. 
The U.S. government spent many millions of dollars on this 
project. What is the present state of this plant? Consult your 
teacher of history or economics or write to your Congressman, 
Senator, or the Tennessee Valley Authority at Knoxville, Ten- 
nessee. What do you think should be done? Prepare a debate 
or discussion of this question. 

2. Write a report comparing the changes brought about by 
the introduction of gunpowder into warfare with the changes 
that the A- and H-bombs may produce. 

3. If available at this time, obtain a fresh sample of some 
legume such as peas, beans, alfalfa, or soybeans. Examine its 
roots for nodules of nitrogen-fixing bacteria. Exhibit to class 
and explain its function. 




. . . In 1858 the atomic theory of 
Da I ton was just 50 years old. Stu- 
dents at this time were generally un- 
familiar with the word molecule, for 
chemists spoke as complacently 
about an atom of water as about an 
atom of oxygen. For the most part, 
also, they had never heard of Amedeo 
Avogadro. William Tilden, 1921 

The battle over the molecules of Avogadro. In 1860, chemical 
science was in a turmoil caused by a misunderstanding of the terms 
atom and molecule. Chemists had spoken of "atoms of water," which 
is a compound, in the same way in which they mentioned atoms of 
hydrogen, which is an element. Some used the term compound 
atoms. So great was the confusion that finally a congress of chem- 
ists was called at Karlsruhe to decide when to use the term atom and 
when to use the term molecule, which up to that time had been used 

Among the brilliant men who gathered at Karlsruhe from all parts 
of the world was a bearded young Italian, Cannizzaro (kan-net- 
sa'ro) . He had come to champion the use of the term molecule in 
the sense that it was used in 1811 by Avogadro (a-vo-ga'dro) , a mod- 
est professor of chemistry. According to Avogadro, a molecule is the 
smallest part of either an element or a compound zvhich has the prop- 
erties of that substance. 

Avogadro had reached this new meaning of a molecule from his 
study of the behavior of gases. He had died four years before the con- 
gress was held, but Cannizzaro championed his ideas so successfully 
that they were finally accepted by the congress. 




What Boyle and Charles discovered about the behavior of gases. 
In order to understand the significance of Avogadro's contribution 
to chemistry, it is necessary to trace the story of the study of gases 
after 1660. In that year Robert Boyle, revered by Englishmen as 
the father of modern chemistry, discovered that if the pressure on 
a gas is doubled for example, increased from 15 pounds to 30 
pounds per square inch its volume is decreased one-half. Further- 
more, he found that this relationship between pressure and volume 
does not depend upon the nature of the gas; it is true for all gases. 
Thus we have Boyle's law: The volume of a gas varies inversely as 
the pressure exerted upon it if the temperature remains constant. 

In 1785, Charles, a French scientist, noted that under constant 
pressure the volume of a gas increases ^^ of its volume at 0C. for 
each centigrade degree of rise in temperature; in other words, if the 
absolute temperature (see page 644) of a gas increases from 273 to 
546 (2 X 273) its volume doubles also. This he found true of any 
gas. Thus we have Charles' law: The volume of a gas varies directly 
as the absolute (A) temperature, if the pressure on the gas remains 

Although gases act according to Boyle's and Charles' laws under 
ordinary temperatures, they do not do so at very high pressures or 
very low temperatures. However, a discussion of the way in which 
gases act at low temperatures and under high pressures is beyond the 
scope of an introductory course in chemistry. 

Gay-Lussac finds the law of combining volumes of gases. The 
next important advance came 23 years later. Gay-Lussac, the French- 
man who collected air samples in a balloon high over Paris, had 
long been interested in the study of gas volumes. He knew that one 
volume of nitrogen unites with one volume of oxygen, forming two 
volumes of nitric oxide. Besides, when he repeated the experiments 
of Cavendish and Lavoisier, he found that two volumes ot hydrogen 
unite with one volume of oxygen, forming two volumes of water va- 
por. This was not a new discovery, but Gay-Lussac suspected that 
"other gases might also combine in simple ratios." 

plunger I 



Fig. 63. A demonstration of 
the principles of Boyle's 


Resuming his researches, Gay-Lussac discovered that one .volume 
of hydrogen chloride gas when brought in contact with one volume 
of ammonia gas yielded a white powder, with no residue of either 
gas. The two gases had combined volume for volume. Furthermore, 
he had read that one volume of nitrogen combined with exactly 
three volumes of hydrogen, forming exactly two volumes of ammo- 
nia gas. This was an arithmetical simplicity of remarkable signifi- 
cance. Fractions of volumes of gases were not involved. 

1 vol. nitrogen + 1 vol. oxygen - 2 vol. nitric oxide 
1 vol. oxygen + 2 vol. hydrogen 2 vol. water vapor 
1 vol. HC1 + 1 vol. NH 3 - NH 4 C1 (a solid) 

1 vol. nitrogen -f 3 vol. hydrogen * 2 vol. ammonia gas 

On the last day of the year 1808, Gay-Lussac formulated from these 
observations a law, which bears his name. Gay-Lussac's law states 
that the relation between the combining volumes of gases and the 
volumes of their products (if they, too, are gases) may be expressed 
in small whole numbers. Why this regularity? On the basis of Dai- 
ton's atomic theory, chemists could not explain this law. 

When Dal ton was faced with this fact, he refused to accept Gay- 
Lussac's law. "The truth is," Dalton maintained, "that gases do not 
unite in equal or exact proportions in any one instance. When they 
appear to do so, it is owing to the inaccuracy of our experiments." 
Later, however, after further experimentation and study, Dalton 
accepted Gay-Lussac's law. His acceptance of Gay-Lussac's law, when 
experimental evidence pointed toward its accuracy, reveals Dalton 
as a true scientist. 

Is the sum of 2 and 1 always 3? The law of Gay-Lussac in par- 
ticular, and the laws of Boyle and Charles to a lesser degree, sug- 
gested a number of interesting problems to Avogadro's inquiring 
mind. Why, for example, is the behavior of gases so uniform under 
changing temperature, while the behavior of solids and liquids is 
so variable? Why do gases combine in simple ratios by volume? Fur- 
thermore, why, with respect to gases, is not the sum of two and one 
always three? For example, why do two volumes of hydrogen unite 
with one volume of oxygen, making two volumes of water vapor; and, 
similarly, three volumes of hydrogen combine with one volume of 
nitrogen, making two volumes of ammonia gas? 

Avogadro continues his work with gases. Avogadro tried to answer 
these questions. In reference to the last question, perhaps he thought: 
it might be that equal volumes of gases contain the same number of 


But according to this idea, one volume of oxygen ought to combine 
with one volume of nitrogen, forming one volume of nitric oxide: 

N + O - NO 

whereas, according to actual experiment, two volumes of nitric oxide 
are formed. Something was wrong. But wait! Suppose the molecule 
of nitrogen gas contains two atoms instead of one, that is, is N 2 and 
not N, and similarly the oxygen molecule is O 2 and not O, what then? 
According to this idea, the equation would be: 

N 2 + O 2 -> 2NO 

and the conditions necessary for the formation of two volumes of NO 
would be fulfilled. 

How would this idea work out in other cases, for example, in the 
formation of water? If the molecule of hydrogen contains two atoms, 
like the molecule of oxygen, we should have the equation, 

2H 2 + O 2 -> 2H 2 O 

or stated in other words, two volumes of hydrogen unite with 1 vol- 
ume of oxygen, forming two volumes of water vapor. This agrees 
with actual measurements of the volumes. 

Apparently Avogadro was on the right track. It remained for him 
only to test his hypothesis further by means of other gas combina- 
tions to be able to show that, assuming the molecules of elementary 
gases to be composed of two atoms each, the volumes corresponded 
with the equation as he had calculated. 

This he actually did and finally was able to establish the accuracy 
of his hypothesis, that all gases behave alike, because equal volumes 
of all gases under the same conditions of temperature and pressure 
are composed of the same number of molecules (Avogadro's hypoth- 
esis) . To this professor from Turin, elementary gases, such as hydro- 
gen, oxygen, nitrogen, and chlorine, consist normally of molecules 
each composed of two atoms instead of one, as Dalton and the rest 
of the world had supposed. Incidentally, Avogadro's hypothesis rec- 
onciled the atomic theory of Dalton with Gay-Lussac's law. How? 

Of what value to chemistry was Avogadro's hypothesis? What 
evidence did Avogadro have to back up so bold an hypothesis? He 
could not verify it experimentally. No balance was sensitive enough 
to weigh a molecule. It would take billions of these tiny particles 
to turn the scales of even the most sensitive balance. He surely had 
not looked into the molecules of matter and detected the twin 


arrangement of atoms, for it would take many millions of mole- 
cules placed side by side to make a line one inch long. Only in re- 
cent years have methods been developed which can make such a 
tiny particle visible. 

The only evidence Avogadro had was that of clear, accurate rea- 
soning and his own creative imagination. However, this evidence 
was strong enough to clear the air and allow chemistry to advance. 
The particles of elementary gases were henceforth considered to be 
diatomic, that is, composed of two atoms to the molecule. (Later, 
by other methods, the inert gases of the atmosphere were shown to 
contain only one atom to the molecule.) Atomic weights and molecu- 
lar weights were thus clearly differentiated. New methods were made 
possible for determining the molecular weights of gases and, from 
these, their atomic weights also. 

How Avogadro's hypothesis was actually verified. Since the time 
of Avogadro, new apparatus and new methods have been devised for 
verifying his prophetic statement. A number of scientists, Millikan 
and Perrin (pe-raN') among them, determined by experiment the 
number of molecules in a given volume of gas. They found that the 
number of molecules in two grams of hydrogen gas (a gram-molecu- 
lar weight) , for example, is approximately 602,000,000,000,000,000,- 
000,000 (602 sextillion) . This number, usually written 6.02 X 10 23 , 
is now called Avogadro's number. Approximately this number of 
molecules is known to be present in equal volumes (22.4 liters) of 
all gases and vapors under the same conditions of temperature and 
pressure. This is no idle guess. 

Perrin and Millikan, both Nobel prize winners in physics, main- 
tained that we can count the number of molecules in a small volume 
of a gas with as much accuracy as we can determine the population 
of a city such as New York. Avogadro's hypothesis has now taken its 
place as one of the laws of chemistry. 

Then came another remarkable verification. Irving Langmuir 
(lang'mur) , another Nobel laureate in chemistry, succeeded in break- 
ing up the molecules of hydrogen gas. As a result of his experiments 
he found that hydrogen gas is made up of molecules each of which 
consists of two atoms. Langmuir made use of this discovery when he 
invented the atomic-hydrogen torch. 

Principle of the atomic-hydrogen torch. In an atomic-hydrogen 
torch, hydrogen gas is passed through an electric arc produced be- 
tween electrodes made of wolfram. The heat of the electric arc splits 
the hydrogen molecule into hydrogen atoms. Immediately after pass- 
ing through the arc, the atoms reunite, forming hydrogen molecules, 

General Electric Company 

Repairing worn parts of vacuum cleaners with an atomic-hydrogen arc welder. 

which are oxidized, forming water. Atomic hydrogen cannot be 
stored. ^ 

H 2 *=> H + H 

All the energy absorbed from the electric arc in splitting the mole- 
cule is liberated when the atoms of hydrogen reunite. This energy, 
as heat, added to the heat normally generated when hydrogen burns, 
produces a temperature between 4000C. and 5000C. (The oxy- 
acetylene torch gives a temperature of about 3300 C.) 

The atomic-hydrogen torch is used for cutting and welding metals. 
It has the advantage of protecting the object against oxidation, since 
the jet of burning hydrogen is always surrounded by hydrogen, a 
reducing agent. 

Proof that the molecule of hydrogen contains two atoms. The- 
oretically, we can prove the formula of hydrogen gas to be H, as 
follows: (1) From experiments, we know that one volume of hydro- 
gen unites with one volume of chlorine, yielding two volumes of hy- 
drogen chloride gas. (2) According to Avogadro's law, equal vol- 
umes of all gases contain the same number of molecules. Conversely, 
equal numbers of molecules of gases occupy equal volumes. Hence, 



one molecule of hydrogen and one molecule of chlorine occupy ,equal 
volumes, and two molecules of HC1 gas occupy twice this volume. 
We may represent this graphically as follows: 

1 volume 4- 1 volume +~ 2 volumes (by experiment) 

1 molecule 4- 1 molecule ^ \2 molecules/- --(by Avogadro's 

of hydrogen \ of HCI / law) 

must contain 1 must contain 

' +~ at leasF*2"atoms of hydrogen 

(3) One of the HCI molecules must contain at least one atom of hy- 
drogen, since fractions of atoms do not exist. (4) Since we have two 
HCI molecules, we must have at least two atoms of hydrogen which 
can have come from only the one molecule of hydrogen. 


1. We know that 1 vol. of nitrogen unites with 1 vol. of 
oxygen, forming 2 vol. of nitric oxide. Prove that the mole- 
cule of nitrogen contains two atoms. 

2. One vol. of hydrogen unites with 1 vol. of bromine vapor, 
forming 2 vol. of hydrogen bromide gas. Show that the mole- 
cule of bromine vapor contains two atoms. 

3. One vol. of oxygen unites with 2 vol. of hydrogen, form- 
ing 2 vol. of water vapor. Show that the formula for oxygen 
gas is O 2 . 

4. Prove the structure of the hydrogen molecule from the 
fact that 1 vol. of nitrogen unites with 3 vol. of hydrogen, 
forming 2 vol. of ammonia gas. 

The gram-molecular volume of a gas or vapor. You have learned 
that a chemical formula may represent one molecule of a compound; 
one molecular weight of a compound; and also one gram-molecular 
weight, or mole, ot a compound. For example, CO 2 may stand for 
one molecule of carbon dioxide, for the molecular weight of car- 
bon dioxide (44) , or for one mole of carbon dioxide (44 grams) . If 
a coefficient appears in front of a formula, it represents a definite 
number of molecules, molecular weights, or moles. Thus 2CO 2 stands 
for two molecules of the gas, two molecular weights (88) , or two 
moles (88 grams) . A formula has an additional meaning which is 
important in many chemical calculations. 


In dealing with a gas or vapor, it is often necessary to know the 
volume that a quantity of it occupies. The unit of measurement of 
gas volumes is the volume occupied by one mole (abbreviated M) . 
This is called the gram-molecular volume (V) . Study has shown that 
the gram-molecular volume is the same for all gases. Under standard 
conditions of temperature and pressure, one M of any gas or vapor oc- 
cupies 22.4 liters. This may be demonstrated by the experimental 
process of weighing a given volume of any gas. For example, one 
liter of hydrogen weighs 0.08987 gram. Therefore 22.4 liters weigh 
2.016 grams, which is the gram-molecular weight (mole) of hydrogen. 

Since one mole of any gas occupies 22.4 liters, we may use the 
formula of the gas to represent its gram-molecular volume. For exam- 
ple, CO 2 , which represents one mole of carbon dioxide, also repre- 
sents one gram-molecular volume (V) of carbon dioxide; NH 3 rep- 
resents one M of ammonia and also one V of ammonia. In each case, 
V = 22.4 liters. If a coefficient appears in front of the formula of a 
gas, it represents the number of gram-molecular volumes. Thus 2CO 2 
stands for 2V (44.8 liters) ; 2NH 3 also stands for 2V. How many liters 
of gas would be represented by 4CO,? by 



In these problems, the weight of one substance is given and 
the volume of another is to be found. Or the volume of one 
is given, and the weight of another is to be found. The pro- 
cedure is the same in both cases. Standard conditions of tem- 
perature and pressure (S.T.P.) are assumed. 

EXAMPLE: How many liters of nitric oxide can be prepared 
by action of sufficient dilute nitric acid on 127.2 g. of copper? 


1. Write the balanced equation. 

3Cu + 8HNO 3 - 3Cu(NO 3 ) 2 4- 4H 2 O -f 2NO 

2. Write the given weight over its formula and x 1. over the 
formula whose volume is to be found. Cross out all other 

127.2 g. x 1. 

3Cu +-SHN03 ->4fN0i); +-4HsO-+ 2NO 


3. Under the formula whose weight is given, write its molec- 
ular weight. Under the formula whose volume is to be found, 
write its gram-molecular volume (V) , not its molecular weight. 

127.2 g. x 1. 

3Cu - 2NO 

3(63.6) - 190.8 2V - 2(22.4) = 44.8 

4. Write the mathematical equation based on the relation- 

wt. of substance used vol. of substance formed 

mol. wt. of substance used V of substance formed 

Solve for x. 127.2 _ xl. 

190.8 " 44.8 
190.8* = 127.2(44.8) 

x = 29.9 liters of NO 

The same general procedure is followed in finding the 
weight of one substance when the volume of another is given 
except that x represents the unknown weight rather than the 
unknown volume. Use an equation based on the same rela- 
tionship for your final solution. 


1. What volume of H 2 may be obtained by the electrolysis 
of 90 g. of H 2 O? 

2. How many liters of NH 3 can be formed by the action of 
33 g. of (NHJ 2 SO 4 on sufficient Ca (OH) 2 ? 

3. How much NaCl is needed to produce 112 1. of HC1 gas? 

Nad + H 2 S0 4 -> NaHSO 4 + HC1 

4. What weight of H 2 O must be decomposed to produce 
224 ml. of O 2 ? 

5. A manufacturer requires 10,000 1. of N 2 O. What weight 
of NH 4 NO 3 must be decomposed? 


This type of problem involves finding the volume of one gas 
or vapor when jhe volume of another is known. As we have 
learned, the coefficient before the formula of a gas represents 
the number of gram-molecular volumes of the gas. Since we 
are dealing only with volumes, weights are disregarded. Only 
the volumes as represented by the coefficients are considered. 



EXAMPLE: How many liters of carbon dioxide are formed dur- 
ing complete combustion of seven liters of benzene, C H ? 

' 66 


1. Write the balanced equation. 

2C 6 H 6 + 15O 2 - 12CO 2 + 6H 2 O 

2. Write the given volume above its formula. Write x L 
above the formula whose volume is to be found. Cross out all 
other formulas. , " 

7 liters x I. 

3. Write the number of gram-molecular volumes (shown by 
the coefficients) under the respective formulas. 

7 liters x 1. 
2C 6 H 6 - 12C0 2 

2 12 

4. Write out the mathematical equation based on the re- 

_ vol. of substance used vol. of substance formed 

coefficient of substance used coefficient of substance formed 

Solve for x. 7 _ x 1. 

2 "12" 
2* = 84 
x - 42 liters of CO 2 


1. 50 1. of H 2 react completely with C1 2 . What volume of 
HC1 gas is formed? 

2. What volume of H 2 is necessary to unite with 5 1. of O 2 
without leaving any O 2 in excess? 

3. What volume of NH 8 can be made from 5000 1. of pure 
N 2 ? 

4. How many liters of O 2 will be used during the complete 
combustion of 500 ml. of methane, CH 4 ? 

GH 4 + 20 2 -> G0 2 + 2H 2 

5. What volume of oxygen will convert 50 ml. of NO into 
nitrogen dioxide, NO 2 ? 



Jaffe, Bernard. Chemical Calculations. World Book Co., 
Yonkers, N.Y., 1947. Additional problems of the types dis- 
cussed in this chapter, together with a more detailed account 
of methods of determining atomic weights and molecular 

Jaffe, Bernard. Crucibles: The Story of Chemistry f pp. 157- 
174. Simon and Schuster, New York, 1948. "The Spirit of a 
Dead Man Leads a Battle" tells the story of Avogadro. 

Perrin, Jean B. Atoms. D. Van Nostrand Co., New York, 
1923. In this book, Perrin, who won the Nobel prize for his 
work on the molecule, describes his verification of Avogadro's 


1. A molecule is the smallest part of either an element or a 
compound which has the properties of that substance. 

2. Gay-Lussac's law states that the relation between the 
combining volumes of gases and the volumes of their products, 
if gaseous, may be expressed in small whole numbers. 

3. Boyle's law states that the volume of a gas varies in- 
versely as the pressure exerted upon it if the temperature re- 
mains constant. 

V 1 /V 2 = P 2 /P! 

4. Charles' law states that the volume of a gas varies directly 
as the absolute temperature if the pressure on the gas remains 

V 1 /V 2 = Ti/T, 

5. Avogadro's law states that equal volumes of all gases 
under the same conditions of temperature and pressure are 
composed of the same number of molecules. 

6. Avogadro's law is valuable because (1) it shows that the 
molecules of certain elementary gases, among them hydrogen, 
oxygen, nitrogen, and chlorine, consist of two atoms; (2) it 
makes possible the determination of the molecular weights of 
gases; (3) it makes possible the determination of the atomic 
weights of gaseous elements; and (4) it shows the relation- 
ships among several apparently conflicting facts concerning the 
actions of gases. 

7. Avogadro's hypothesis was verified by the work of other 
scientists and today is a chemical law. The actual number of 


molecules in the gram-molecular weight of a gas was deter- 
mined by experiment. This number, called Avogadro's number, 
is the same for all gases. It is 6.02 X 1Q28 - 

8. The gram-molecular volume of a gas is the volume occu- 
pied by its gram-molecular weight. Under standard conditions, 
it is 22.4 liters. 


Group A 

1. In 1860, what was the condition of chemical usage with 
respect to the terms atom and molecule? 

2. What is the difference between an atom and a molecule? 

3. When and by whom was the term molecule first clearly 

4. What did Cannizzaro do to establish the meaning of 

5. (a) State Gay-Lussac's law and (b) give two illustrations 

ofit ' ' 


6. The volume of a gas changes from 10 to 5 1. when the 
pressure on it changes from 1 to 2 atm. What law does this 

7. The volume of a gas changes from 4 1. to 2 1. when its 
temperature changes from 500A. to 250A. State the law 

8. State Avogadro's law. 

9. Upon what facts did Avogadro base his hypothesis? 
10. State two ways in which Avogadro's law is valuable. 

1 1 . What two scientists verified Avogadro's hypothesis? 

12. What is a gram-molecular volume of a gas? 

13. Outline the method used in working a weight- volume 

14. Outline the method used in working a straight- volume 

15. If 15 1. of N 2 are needed to unite with O 2 in forming 
NO, what volume of O 2 will be used? 

t . . . 


16. Assume that air contains 20 percent O 2 by volume. What 
volume of air will be needed in forming 100 ml. of O 3 ? 

17. What volume of air will be needed for the complete 
combustion of 750 ml. of acetylene, C 2 H 2 ? 

18. CO passed over warm Ca (OH) 2 reacts as follows: 

CO + Ca(OH) 2 - CaCO 3 + H 2 | 

How does the volume of CO compare with that of the H 2 ? 

19. What weight of carbon is in 44.8 1. of CO? 

20. What volume of NO 2 will be formed by the complete 
reaction of 100.5 g. of Hg with concentrated HNO 3 ? 

Hg + 4HN0 3 -> Hg(N0 3 ) 2 + 2H 2 O + 2NO 2 

21. HC1 gas was bubbled through a solution of NaOH. As 
a result, 468 g. of NaCl were formed. What volume of the 
HC1 used actually combined with the base? 

Group B 

22. From the experimental fact that 3 vol. of O 2 change into 
2 vol. of O 3 when an electric discharge is passed through moist 
O 2 , prove that the molecule of ozone contains three atoms. 

23. Prove nitrogen molecule contains at least two atoms. 

24. Which is more economical to use in the preparation of 
NH 3 , (NH 4 ) 2 SO ? at $8.75 per 100 Ib. or NH^Cl at 12^ per lb.? 

25. (a) Explain the operation of the atomic-hydrogen torch, 
(b) Account for the extreme heat obtained. 

26. What is meant by absolute temperature? 

27. Prove by calculation that the ounce-molecular-volume 
of any gas equals 22.4 cu. ft. 


1. In England, Robert Boyle is considered the father of 
modern chemistry. In France, Lavoisier is called the father of 
modern chemistry. Can you suggest reasons for this difference 
of opinion among French and English scientists? Is it justified? 
Does it illustrate scientific open-mindedness? 

2. Suggest possible reasons for the neglect of Avogadro's 
hypothesis from 1811 to 1860. Can you cite other scientific 
work which remained unrecognized for a long time? 

3. Construct a cardboard or wooden box to represent the 
gram-molecular volume of any gas at S.T.P. Exhibit it to your 
class alongside of a quart bottle or carton of milk. 





. . . Sulfur has been taken intermit- 
tently from Popocatepetl since the 
time of the ancient Aztecs, who used 
it for medicinal purposes. Two of 
Cortez's soldiers, in the sixteenth 
century, climbed to the crater and 
obtained sulfur for the purpose of 
manufacturing gunpowder. Science 
News Letter 

An American pharmacist creates a new industry. The discovery 
of petroleum in Pennsylvania in 1859 led at once to a wide search 
for other stores of oil. Only six years later, oil prospectors stumbled 
upon huge deposits of sulfur, a yellow, brittle, lustrous solid known 
since ancient times. These deposits were about 500 feet below the 
surface in Louisiana not far from the Gulf near the Texas line. Here 
ages ago a vast geyser had spouted, leaving the sulfur within and 
about its crater. The sulfur was covered with strata of clay, limestone, 
and, worst of all, gas and quicksand. It was impossible to sink shafts 
to reach the deposits in order to dig the sulfur out as coal is mined. 
Many companies were formed to exploit these deposits, but because 
of the many difficulties, they all failed. 

In 1891 Herman Frasch (frash) heard about this sulfur. He had 
come from Germany 25 years before. Leaving high school at 16, he 
had been apprenticed to a druggist and then had left for the United 
States, where he opened his own drugstore in Philadelphia. Chem- 
istry fascinated him, and in the back of his drugstore he carried on 
many researches on petroleum products. Later he sold his store and 
devoted all his time to chemical engineering. He tackled the problem 
of extracting sulfur from the Louisiana deposits. 



u alr 


melted sulfur 

' '' : /' :' V"'.\ : ''.'" anhydrite " / ,'v 

Fig. 64. The Frasch process for extracting sulfur from deposits deep in the earth. 

When Frasch described to some of his friends the entirely new 
process that he had devised, they thought it impossible. One promi- 
nent person challenged him in no uncertain terms. He said that he 
would swallow every ounce of sulfur Frasch extracted by his queer 
process of pumping a solid out of the earth. But Frasch kept on im- 
proving his method and ultimately he succeeded in founding a new 
American industry based on his process. 

The Frasch process of extracting sulfur. Frasch' s plan was to sink 
a well by means of an oil-drilling rig, and lower three concentric 
pipes (inside a casing) down to the sulfur. Through the outermost 
6-inch pipe, superheated water was to be pumped to melt the sulfur. 
Through the innermost one-inch pipe, compressed air was to be 
forced down to the sulfur deposit to make the melted sulfur frothy 
and light. The result was to be a flood of molten, foamy sulfur gush- 
ing under pressure from the three-inch pipe between the other two. 

Frasch was visibly nervous when he gave the order to start the 
first pump. More and more slowly went the engine with its increas- 
ing load until the man at the throttle sang out at the top of his voice, 
"She's a-pumping." A liquid appeared at the mouth of the three-inch 
pipe. Frasch wiped the liquid off a polished piston rod, and gazed 
upon the first crystals of sulfur extracted by his ingenious process. 
Then came a steady stream of golden liquid, which, in 15 minutes, 
filled every one of the 40 barrels available. Still the molten sulfur 
kept pouring. Embankments were quickly thrown up and lined with 
boards to hold the sulfur as it solidified. 

This is the method of mining sulfur in use today in Texas and 
Louisiana. These deposits now supply much of the world's needs. In 
this way, mountains of sulfur, 99^ percent pure, are extracted and 
stand ready to be dynamited into pieces small enough to be loaded 
and shipped to many parts of the world. 



Sulfur is obtained in Sicily by a different process. Before the time 
of Frasch, a group of English financiers had been marketing the rich 
sulfur deposits of the volcanic region of Sicily. Here the sulfur is 
found mixed with clay and limestone, from which it is separated by 
melting this ore, allowing the free sulfur to flow away from the im- 
purities. The sulfur is then boiled and is changed to a powder, called 
(loiuers of sulfur, by chilling the vapors. 

The English monopoly had been supplying more than 90 percent 
of the world's sulfur. In 1904, when American output reached the 
point where a single well could supply 400 tons of sulfur daily, the 
English company went out of business. To prevent the unemploy- 
ment of hundreds of workers, the Italian government took control. 
Frasch aided gladly in stabili/ing the Sicilian sulfur industry. 

Crude free sulfur accounts for about only 40 percent of the world's 
sulfur production. Iron, zinc, and copper sulficles, natural gas, and 
industrial gases supply the rest. The United States today leads the 
world in the production of sulfur. Many countries, including Italy 
and Mexico, are stimulating their own production of sulfur by sub- 
sidies, tariff laws, bounties, and price guarantees. 

Physical properties of sulfur. Sulfur is a pale yellow, soft, brittle 
solid (plastic, or pliable, in the case of amorphous sulfur) with a 
very faint odor and no marked taste. It is practically insoluble in 
water, more soluble in carbon tetrachloride, CC1 4 , and very soluble 
in carbon disulfide, CS... 

Sulfur is a poor conductor of heat. It melts at about 114C., 
forming a pale yellow liquid, which on further heating darkens and 
thickens, becoming- almost black at 235 C. At a still higher tempera- 
ture, it becomes thinner again, and finally changes to a yellow vapor 
at 445C. Sulfur is mentioned in the Bible as brimstone. 

Texas Gulf Sulfur Cohipany 

Liquid sulfur flows into the 
storage vat in which it will 
solidify. The pipeline leads 
to the sulfur well. 



Allotropic forms of sulfur. Oxygen, as you have learned, occurs in 
two allotropic forms, O 2 and O 3 , but sulfur occurs in two common 
crystalline forms (rhombic and prismatic) , and one noncrystalline 
or amorphous form. Each has different properties caused by differ- 
ences in the arrangement of their atoms. 

1) Rhombic sulfur. Sulfur is found in nature in the form of 
rhombic crystals, the most stable form under normal conditions. 
Their molecules consist of rings of eight atoms of sulfur (S s ) , and 
their density is about two. Crystals of this form of sulfur may be 
prepared by dissolving sulfur in carbon disulftde and allowing the 
solvent to vaporize slowly. The residue consists of perfect crystals 
having the shape of two pyramids with their bases joined. Roll sulfur, 
made by cooling sulfur in cylindrical molds, is almost entirely 

2) Prismatic sulfur. When sulhir is heated until it just melts and 
is then allowed to cool slowly, it forms long needle-shaped crystals 
whose density is somewhat less than that of rhombic sulfur. 

3) Amorphous sulfur. When boiling sulfur is suddenly cooled by 
pouring it into cold water, the rings are broken and are replaced by 
long chains of sulfur atoms which become entangled and can be 
stretched like rubber. It is amber in color and, unlike the other 
two forms, is insoluble in carbon disulfide. Amorphous sulfur 
changes gradually into rhombic at room temperature. For simplicity, 
all forms are designated S. Flowers of sulfur is a powdered mixture 
composed of rhombic and some of the plastic. Between 96 C. and 
1 14C. the most stable form of sulfur is the prismatic. 

Chemical properties of sulfur. The atomic weight of sulfur is 
32. Its atomic number is 16, and it has, therefore, six electrons in its 
third, or outermost, ring. This makes it a borrower of electrons. It 
is a nonmetal, fairly active, and has, under ordinary conditions, a 
valence of minus two. Therefore, sulfur closely resembles oxygen in 
its chemical properties. 

Fig. 65. Allotropic form, of sulfur. melfed Sulfur " 



Exterior of a sulfur storage vat. 
The solidified sulfur will be 
blasted into small pieces for 
shipment in railroad cars. 

Texas Gulf Sulfur Company 

Like oxygen, sulfur unites with most metals. The compounds of 
metals and sulfur are called sulfides. For example, when sulfur is 
heated with iron powder, iron sulfide is formed. 

- - Fe-fS-^FeS 

Hot copper burns in sulfur vapor, forming cuprous sulfide. 

. /- V^ 2Cu+S-+Cu 2 S 

When sulfur is mixed with zinc dust and ignited, the chemical 
reaction is so vigorous that a great deal of light and heat are liber- 
ated and dense clouds of zinc sulfide, which settle out as a powder, 
are formed. 

Zn + S-^ ZnS 

Although sulfur is a nonmetal, it is less nonmetallic than oxygen. 
It therefore combines with oxygen, exhibiting a positive valence of 
either four or six. Sulfur burns in air with a pale blue flame, form- 
ing sulfur dioxide, SO.,. A small amount of sulfur trioxide, SO 3 , is 
formed later by further oxidation of the SO 2 . 

2SO 2 + O 2 

> SO 2 (sulfur dioxide) 

> 2SO 3 (sulfur trioxide) 

The electron structure of sulfur and oxygen are shown below. 
Fig. 66. The structure of the oxygen atom (left) and the sulfur atom. 





Sulfur also shows more metallic properties than chlorine by unit- 
ing with the latter, forming sulfur dichloride, SCI.,, a brownish red 
liquid with a pungent odor used in chlorinating, and also sulfur 
monochloride, a heavy, amber-colored unstable liquid with an irri- 
tating odor, used in vulcanizing rubber. 

2S + C1 2 -> S 2 C1 2 (sulfur monochloride) 

When sulfur vapor is passed over carbon heated in an electric fur- 
nace, the two elements combine, forming carbon disulfide, CS 2 . 

C + 2S - CS 2 

Carbon disulfide is a heavy, colorless liquid with a pleasant odor. 
Generally, however, it has a disagreeable odor caused by impurities. 
It is very combustible. Its chief use is as a solvent for sulfur, gums, 
rubber, fats, and waxes. It has been used also as a poison in exter- 
minating ants and other insects, rats, and prairie dogs. 

Rubber tires depend on sulfur. The normal annual consumption 
of sulfur in the United States is more than three million tons. It is 
one of the fundamental industrial elements. By far the greatest 
quantities of sulfur are used in the manufacture of sulfur ic acid, 
one of the most widely used of the heavy chemicals. The production 
and uses of sulfuric acid are discussed in Chapter 21. 

Great quantities of sulfur are used in vulcanizing rubber. By this 
process, soft, sticky, perishable, natural rubber is changed to a harder, 
less plastic, more durable rubber. Vulcani/ed rubber is used chiefly 
in making automobile tires but also in making thousands of other 
rubber articles. 

Tfie Firestone Tire and Rubber Company 

In a modern tire factory, a worlc- 
man removes a finished tire 
om a steam-heated vulcani- 
zing mold. 


When the American, Charles Goodyear, in 1839 while working in 
his kitchen in Woburn, Massachusetts, accidentally dropped a piece 
of rubber mixed with sulfur on a hot stove, he discovered the proc- 
ess of vulcanization (derived from Vulcan, the Roman god of fire) . 
A new and highly important industry was made possible. To shorten 
the time required for vulcanizing rubber, a catalyst, or accelerator, 
such as zinc oxide, is added to the mixture of rubber and sulfur. 

In making an article of rubber, the washed raw rubber is thor- 
oughly mixed with various chemicals that determine the properties 
of the finished product. Among these substances are sulfur, carbon 
black, lead oxide (PbO) , zinc oxide, and carbonates of magnesium 
and calcium. Different combinations of these and other substances 
in varying quantities may be used in accordance with the properties 
desired in the finished product. The rubber is then rolled into sheets 
of the desired thickness or placed in molds of the desired shape. 
While in the mold, the rubber is heated and vulcanization occurs. 

Sulfur is used in controlling fungi and insects. A very effective 
liquid for destroying fungus growths and fruit tree, shrub, and vine 
pests is a deep orange-red lime-sulfur spray made by boiling sulfur 
with calcium hydroxide. The spray is both a fungicide, a substance 
that kills fungi and molds, and an insecticide, a substance that kills 
insects. Dusting with very finely powdered sulfur is effective against 
rose diseases, mildew, and black spot. However, ordinary flowers of 
sulfur is not fine enough to be of much value as a dusting agent, and 
even some finely ground commercial dusting powders are too coarse. 

Colloidal sulfur and wettable sulfur, the first a very highly dis- 
persed sulfur in water and the second sulfur so treated that it dis- 
perses on contact with water, are both used in making mild sprays 
that are particularly useful for the summer spraying of roses and for 
the control of mildew and true rust diseases of other plants. 

Sulfur is used in medicine. Ointments containing sulfur have 
been used since antiquity to control skin diseases caused by fungi. 
A common ointment of this kind is made by mixing three parts by 
weight of sulfur, 15 parts of white petrolatum, or Vaseline, one part 
of lanolin, or wool fat, and one part of yellow wax. Such an oint- 
ment is effective in killing the very small mites that cause scabies, 
or itch. Sulfur ointments may be of value in the control of infectious 

Other uses of sulfur. Much sulfur goes into the manufacture of 
calcium bisulfite, Ca (HSO 3 ) 2 , which is used in making wood pulp 
for the manufacture of paper. Sulfur is used also in making syn- 
thetic resins, sulfur colors, and gunpowder. Sulfur-lined steel pipes 



are used to transport liquids that are very corrosive to the materials 
of which pipes are normally made. Sulfur cements are used to join 
bricks in floors and walls that are continually subjected to the corro- 
sive effects of acids or alkalies. Great quantities of sulfur are used 
in the making of matches. 

How matches are made. Early matches were dangerous, hard to 
use, and hard to carry. Phosphorus matches often caused painful 
burns and were harmful to the health of the workers who made them. 

The first friction match was invented in 1827 by an English phar- 
macist, John Walker. It contained potassium chlorate (KC1O 3 ) , an 
oxidi/ing agent, and antimony sulfide (Sb,S s ) , a compound with a 
low kindling temperature. The locofoco match was an American 
adaptation of Walker's friction match. During the presidential cam- 
paign of 1840, the Democrats were called locofocos because at one of 
their meetings, they used matchlight when the Whig landlord turned 
off the gas. 

In 1831 white phosphorus was used for the first time in making 
matches. It was a more efficient fire-producer than antimony sulfide, 
but played havoc with the health of the match-factory workers, who 
finally demanded protection. This led in 1906 to an international 
convention which prohibited the further use of white phosphorus 
in making matches because of its poisonous nature. 

The strike-anywhere match in use today also contains both an oxi- 
dizing agent and a compound with a low kindling temperature. The 
head of the paraffin-dipped matchstick contains potassium chlorate 
and phosphorus sulfide (P 4 S 3 ) , a dark solid that is concentrated on 
the tip of the match. Glue is used to bind the chemicals, and ground 
glass or other abrasive is sometimes added as a filler. 

The tip of a safety match is composed of easily combustible anti- 
mony sulfide (Sb.,S 3 ) , and potassium chlorate, which provides addi- 
tional oxygen. The side of the box contains red phosphorus, which 



hydrogen sulfide S 

Fig. 67. Laboratory preparation of 
hydrogen sulflde (left). Structure 
of the hydrogen sulflde molecule 

is nonpoisonous. The material on the tip of the match will not ig- 
nite easily unless it is rubbed on the treated side of the box. 

Disastrous fires are frequently caused by matches that have been 
blown out but still retain glowing tips. To prevent this dangerous 
afterglow, matches are dipped in some such solution as sodium sul- 
late or ammonium phosphate during their manufacture. When 
blown out, they do not leave glowing tips. 

Waterproof matches are treated with a transparent coating with 
a high-kindling temperature. They will light even after being soaked 
in water. 

The manufacturer of matches uses a continuous-process machine, 
which takes pine wood, cuts it, dips the sticks in paraffin, coats them 
with the chemicals needed, dries them, and finally packs them for 
shipment. More than a million matches can be made by one of 
these machines each hour. 

How sulfur and hydrogen unite. Hydrogen has only one electron. 
It therefore needs one electron to complete its only ring; but since 
it shows a strong tendency to lend its one electron, it is said to pos- 
sess metallic properties. Sulfur, requiring two electrons to complete 
its outer ring, will combine with two atoms of hydrogen. This union 
forms the compound, hydrogen sulfide, H,S, as shown in Fig. 67. 
In this compound, sulfur exhibits definitely nonmetallic properties. 
Although it is possible to prepare hydrogen sulfide by direct union 
of sulfur and hydrogen, it is prepared most easily by other methods. 


How hydrogen sulfide is prepared. Hydrogen sulfide is prepared 
both commercially and in the laboratory by the general method of 
preparing an acid, that is, by the action of sulfuric acid on a salt (sul- 
fide) . The sulfide most generally used is ferrous sulfide, FeS, a black 
iron ore. When sulfuric acid is poured on ferrous sulfide in a test 
tube (see Fig. 67) , hydrogen sulfide gas is given off immediately 
without the addition of any external heat. 

FeS -I- H 2 SO 4 -> FeSO 4 + H 2 S | 

Physical properties of hydrogen sulfide gas. Hydrogen sulfide is 
colorless, slightly heavier than air, and fairly soluble in water (one 
volume of water dissolves three volumes of hydrogen sulfide) . Natu- 
ral sulfur waters contain hydrogen sulfide in solution and, upon be- 
ing decomposed, leave a deposit of free sulfur. Easily liquefied, 
hydrogen sulfide is sold in cylinders for laboratory use. Its most char- 
acteristic physical property is its odor, the odor of rotten eggs. In 
fact, it is hydrogen sulfide that gives such eggs their odor. It is caused 
by the decomposition of organic sulfur compounds in the yolks. 

Hydrogen sulfide forms naturally in marshes, oil wells, mines and 
coal piles, manure pits, and sewers. In persons who have endured 
mild exposure to its effects, it produces inflamed thrpat, headache, 
a heavy feeling in the stomach, and dizziness. When breathed in 
large quantities it causes death. Both natural gas and coal gas con- 
tain H 2 S which is removed before they are used as household fuels. 

Chemical properties of hydrogen sulfide gas. When burned in 
sufficient air, hydrogen sulfide gives off a pale blue flame, and water 
vapor and sulfur dioxide are formed. This sulfur dioxide gives burn- 
ing hydrogen sulfide its irritating odor. The equation for the com- 
plete combustion of hydrogen sulfide is: 

2H 2 S + 3O 2 - 2H 2 O + 2S0 2 T 

When hydrogen sulfide is burned in a small amount of air (incom- 
plete combustion) , water is formed as before, but free sulfur is pro- 
duced instead of sulfur dioxide. This sulfur separates out as a yellow 
powder. The fact that it does so probably accounts for the presence 
of sulfur around volcanoes, which emit hydrogen sulfide gas. 

2H 2 S + O 2 - 2H 2 O + 2S T 

Because of the tendency of hydrogen sulfide to unite with oxygen, 
it is a fairly good reducing agent. When hydrogen sulfide acts as a 
reducing agent, the sulfur lends electrons and is oxidized. 



In a water solution, hydrogen sulfide dissociates to some extent. 
Such a solution acts as a weak, unstable acid, sometimes known as 
hydrosulfuric acid., H 2 S. On continued boiling, this acid liberates hy- 
drogen sulfide, leaving pure water. 

Hydrogen sulfide reacts with certain metals and also with the salts 
of certain metals, forming sulfides. The tarnishing of silverware is 
caused by the formation of a brownish-black sulfide of silver, Ag 2 S. 
The blackening of lead paints is caused by the formation of black 
lead sulfide. Lithopone, a white paint base now used widely, consists 
of a mixture of barium sulfate (RaSO 4 ) , and zinc sulfide (ZnS) . 
Lithopoiie paints do not lose their color by the action of H,S. Why? 
Titanium dioxide, TiO 2 , also is used as a white paint base. It is not 
blackened by sulfur compounds and has great covering power. 

Many important sulfides are found in the earth. The salts of 
hydrosulfuric acid form an important class of compounds called 
sulfides. Many of them occur in nature and constitute important ores 
such as iron pyrites., FeS.,; galena, PbS; zinc blende, ZnS; cinnabar, 
HgS; and CuS. Certain colored sulfides are used as mineral pig- 
ments in the coloring of paints. Cadmium sulfide, CdS, for example, 
is a yellow pigment, and zinc sulfide, ZnS, is a white pigment used 
in paints (see illustration following page 382). 

The chief use of hydrogen sulfide. In the analysis of ores and in 
the separation of groups of certain metals from other groups of met- 
als, hydrogen sulfide gas is indispensable. For this reason, a hydro- 
gen sulfide generator is always present in a laboratory for analytical 
chemistry. The sulfides of certain metals such as sodium and calcium 
are soluble in water, while those of other metals such as lead and zinc 
are insoluble. 

By passing hydrogen sulfide into a solution of the soluble salts of 
such metals, the sulfides of certain metals precipitate out and may 
be separated by filtration. Thus, certain metals present in an ore or 

Standard Oil Com puny (N .J .) 

Hydrogen sulfide fumes ris- 
ing from crude oil may 
prove fatal to workmen. 
Hence, protective masks 
are worn by anyone work- 
ing near the oil storage 


alloy may be separated with the aid of H 2 S. Furthermore, since the 
colors of the sulfides of many metals differ, chemists can use these 
differences in color in detecting the presence of these metals. For 
example, zinc sulfide is white, arsenic sulfide is yellow, antimony 
sulfide is orange, and copper sulfide is brownish-black. 

Zn(NO 3 ) 2 + H 2 S -> ZnS I + 2HNO 3 
CuSO 4 + H 2 S -> CuS I + H 2 SO 4 

This difference in color is only one of the very many ingenious 
methods used by analytical chemists in detecting and isolating ele- 
ments present in complex mixtures and compounds. 

The test for the sulfide ion. If, on the addition of sulfuric acid to a 
compound, hydrogen sulfide gas is liberated, the substance tested 
is a sulfide. The hydrogen sulfide liberated is easily detected either 
by its odor or by its ability to turn a silver coin brownish-black, as a 
result of the formation of silver sulfide. 


Fabre, Jean H. The Wonder Book of Chemistry, pp. 345-359. 
Albert &T Charles Boni, New York, 1922. Discusses sulfur and 
includes some simple experiments with this element. 

Waggaman, W. H., and Barr, J. A. "Sulfur for Survival." 
Chemistry, October, 1951, pp. 1-10. An illustrated article on 
sources, extraction, and properties of sulfur. 


Group A 

1. When and by whom was the first successful method of 
obtaining S from the Louisiana deposits invented? 

2. What was the great difficulty that had to be overcome 
before S could be extracted from the Louisiana deposits? 

3. How did the Frasch process affect the Sicilian sulfur 

4. (a) Make a labeled diagram of the Frasch process and 
(b) explain it. 

5. What are two ways in which the Louisiana sulfur de- 
posits differ from the Sicilian sulfur deposits? 

6. Name the allotropic forms of S and tell how each may 
be prepared in the laboratory. 


7. What two elements other than S occur in allotropic 

8. (a) State the chief differences and resemblances of the 
three allotropic forms of S. (b) How could you prove that all 
three forms of S are the same element? 

9. A piece of plastic S is left overnight in the laboratory. 
The next morning yellow brittle S is found. Explain. 

10. (a) How are the two kinds of commercial sulfur, roll 
sulfur and flowers of sulfur, prepared? (b) What allotropic 
forms are in each? 

1 1 . Make a diagram of an atom of S and use it to find the 
valence of S, to describe its chemical activity, and to explain 
why it is a nonmetal. 

12. In what three ways does S resemble O 2 ? 

13. Write balanced equations for the: (a) union of sulfur 
and zinc; (b) complete combustion of hydrogen sulfide; (c) in- 
complete combustion of hydrogen sulfide; (d) burning of sul- 
fur; (e) union of sulfur and chlorine; (f) complete combustion 
of carbon disulfide. 

14. What are four chief uses of S? List in order of im- 

t . . . 

15. (a) What is meant by the vulcanizing of rubber? 

(b) What properties does vulcanizing impart to rubber? 

(c) What American discovered this process? 

16. (a) Of what chemicals is the tip of a strike-anywhere 
match composed? (b) How does a safety match differ from an 
ordinary match? 

17. Hydrogen sulfide present in natural gas is removed be- 
fore the natural gas is sent into the pipelines to be used as fuel. 

18. By a labeled diagram describe the laboratory prepara- 
tion of H 2 S. 

19. What weight of FeS would be needed to prepare 204 Ib. 
of H 2 S? 

20. What volume of gaseous H 2 S could be prepared from 
20 g. of 90 percent pure FeS? 

21. A compound of hydrogen and sulfur has a mol. wt. of 34. 
The percentage of S in the compound is 94.1 percent. Find its 


22. What volume of H 2 S is required to precipitate all the 
CuS from a solution containing 80 g. of CuSO 4 ? 

23. What volume of SO 2 will be formed by the complete 
combustion of 896 ml. of gaseous H 2 S? 

24. Write ionic equations for the following and tell whether 
each reaction goes to completion (refer to Table 10, page 243) : 

a) FeS + H 2 S0 4 - FeSO 4 + H 2 S 

b) Pb(N0 3 ) 2 + H 2 S - PbS + 2HNO 3 

c) ZnSO 4 + (NH 4 ) 2 S -> ZnS + (NH 4 ) 2 SO 4 

d) Na 2 SO 4 + H 2 S - Na 2 S + H 2 SO 4 

25. Compare the physical properties of H 2 S and N 2 O. 

26. Describe what happens when gaseous H 2 S is bubbled 
through water. 

27. Why is H 2 S said to be (a) weak? (b) unstable? 

28. Explain the tarnishing of silverware. 

29. Why is H 2 S solution kept in amber-colored bottles? 

30. Using two equations, show the chief use of H 2 S. 

31. Describe a test for a sulfide. 

32. Copy and complete. Do not write in this book. S occurs 
in our bodies because it is one of the elements found in .... 
The two chief sulfur-producing states in America are Louisi- 
ana and .... A lime-sulfur preparation is used as .... A sulfur 
compound used in exterminating rats is .... A sulfide that is 
soluble in water is .... 

33. Under a rubber band used to keep pieces of silver 
cutlery together, a black mark is found. Explain. 

Group B 

34. How would you tell fool's gold, FeS 2 , from genuine gold? 

35. In the Frasch process, why is S forced out of the three- 
inch pipe rather than out of the six-inch pipe? 

36. The action of H 2 S on a solution of ZnSO 4 is reversible. 
How would you force this reaction to completion? 

37. NH 3 leaks may be detected by burning sulfur candles. 

38. (a) Why can we not use HNO 8 in preparing H 2 S from 
FeS? (b) What is formed if H 2 S is passed into HNO 3 ? 


39. Why do we write CS 2 and not S 2 C? 

40. A mixture of sulfur and molasses was given to children 
as a "spring tonic." What do you think of this practice? 


1. Discovery of the vulcanization of rubber is generally con- 
sidered an accident. Had the accident not happened to Good- 
year, would the thousands of rubber articles in use today never 
have come? Explain your answer fully in terms of scientific 
advance and the needs of society. 

2. If you live in a rural section, talk with as many farmers 
as possible to learn what methods they use to check the losses 
incurred by insect pests and fungus growths. Report your 
findings to your class. 

3. Filter paper dipped in a solution of Pb (C 2 H 3 O 2 ) 2 turns 
black in the presence of H 2 S. Hang small strips of it in the 
basement while the furnace is burning, over the kitchen stove 
or gas range while dinner is being cooked, and in the living 
room. Which of the strips turns blackest first? What do you 




. . . We may fairly judge of the 
commercial prosperity of a country 
from the amount of sulfuric acid it 
consumes. Reflecting upon the im- 
portant influence which the price of 
sulfur exercises . . . we can under- 
stand why the English Government 
should have resolved to war with 
Naples, in order to abolish the siu- 
fur monopoly, which the latter 
power attempted recently to estab- 
lish. Justus yon Liebig, 

Why sulfuric acid is the fundamental chemical today. One of the 

outstanding differences between our society and the society which 
preceded it is the tremendous role that machines play. Automobiles, 
locomotives, steel ships, and airplanes are comparatively new. In the 
manufacture, operation, and upkeep of these and thousands of other 
machines, sulfuric acid is directly or indirectly a prime necessity. If 
the ability to make this acid were lost suddenly, industry would be 
paralyzed. So fundamental is it to our industrial life that its con- 
sumption is a fair index of industrial conditions. 

How sulfur dioxide is prepared for use in industry. About 90 per- 
cent of the sulfur consumed industrially is first burned in air to 
produce sulfur dioxide. Heating sulfide ores is another commercial 
method of preparing SO 2 . When an ore, such as iron pyrites, FeS 2 , is 
heated in air, or roasted, one of the products formed is sulfur dioxide. 

4FeS 2 + 11O 2 - 2Fe 2 O 3 + 8SO 2 T 

Sulfur dioxide occurs in small quantities in the vapors rising from 
active volcanoes and certain sulfur Springs, and in the gases formed 
during the burning of coal. 


Na 2 SO 3 + H 2 SO 4 

Fig. 68. Laboratory prepa- 
ration of sulfur dioxide. 
Why is the excess gas 
passed through water? 

-* H 2 Q 

How sulfur dioxide is prepared in the laboratory. When sodium 
sulfite, Na 2 SO 3 , or any other sulfite is treated with an acid, sulfur 
dioxide is liberated. The gas is collected by the displacement of air, 
as shown in Fig. 68. 

Na 2 SO 3 + H 2 SO 4 - Na 2 SO 4 + H 2 O + SO 2 1 

Physical properties of sulfur dioxide. Sulfur dioxide is colorless, 
has a suffocating odor, is more than twice as heavy as air, and is very 
soluble in water (one volume of water at room temperature dissolves 
40 volumes of the gas) . Under a pressure of only three atmospheres 
(approximately 45 pounds per square inch, or 45 psi) it changes to 
a colorless liquid, which can be transported in steel cylinders or 
shipped like gasoline in tank cars. 

Chemical properties of sulfur dioxide. Sulfur dioxide does not 
burn in air, but it may be made to combine with another atom of 
oxygen, forming sulfur trioxide, SO 3 , by passing a mixture of sulfur 
dioxide and air over a catalyst, such as platinum or vanadium pent- 

2SO 2 + O 2 -* 2SO 3 

This ability of sulfur dioxide to combine with more oxygen makes 
sulfur dioxide a reducing agent. Sulfur dioxide may also act as an 
oxidizing agent. When, for example, it is passed through a solution 
of hydrosulfuric acid, H 2 S, it precipitates sulfur. The sulfur (S++++) 
in the SO 2 takes electrons from the sulfur (S ) of H 2 S and changes 
to free sulfur, S. 

S+++ + O 2 + 2H 2 +S - 2H 2 +O + 3S j 

Hydrosulfuric acid acts here as the reducing agent. These two actions 
of sulfur dioxide illustrate the fact that a substance may be either an 
oxidizing agent or a reducing agent, depending upon the substance 
to be reduced or oxidized. 







check valve- 




H^lpfc^ expansion valve 

freezing tray 

Fig. 69. An electric refrigerator. 
The refrigerant may be sulfur 
dioxide, methyl chloride, Freon, 


electric motor compressor 

When sulfur dioxide dissolves in water, it combines with some of 
the water, forming sulfurous acid. Sulfur dioxide, then, is an acid 

SO 2 + H 2 O > H 2 SO 3 (sulfurous acid) 

Sulfur dioxide has many uses. The most common use of sulfur 
dioxide is in the manufacture of sulfuric acid. It is used also to some 
extent instead of ammonia in mechanical refrigeration. Although 
liquid sulfur dioxide during evaporation absorbs only one-third as 
much heat as ammonia does, it liquefies under a lower pressure and 
hence is used in some household refrigerators. 

Homer in his Odyssey cited the use of burning sulfur in fumiga- 
tion. Today rooms, houses, or warehouses are sometimes fumigated 
by burning sulfur candles or by liberating sulfur dioxide gas from 
cylinders. However, as a fumigant sulfur dioxide has been almost 
wholly supplanted by more effective substances. It is even extremely 
doubtful that fumigation by sulfur candles was ever very effective. 

How sulfurous acid, H,SO 3 , is prepared in the laboratory. In 
the laboratory preparation of sulfur dioxide, SO 2 is not the product 
first formed. Sulfurous acid is the first product of the reaction, but 
it>J impossible to stop the reaction at this point. 

Na 2 SO 3 + H 2 SO 4 - Na 2 SO 4 + H 2 SO 3 

Sulfurous acid is prepared by passing sulfur dioxide gas through 
cold water. 

H 2 O + SO 2 - H 2 SO 3 


Properties of sulfurous acid. Sulfurous acid is a colorless solution 
with a suffocating odor (sulfur dioxide) . It is unstable and decom- 
poses readily into water and sulfur dioxide. By boiling sulfurous 
acicl, all the sulfur dioxide may be driven off. The reaction is re- 
versible, depending upon the temperature. 

H 2 + SO. t H 2 SO 3 

Sulfurous acid is a very weak acid because it dissociates only 
slightly, forming few hydrogen ions. When neutralized with a base, 
it forms salts known as sulfites. 

H 2 SO, + 2NaOH -> Na 2 SO 3 (sodium sulfite) + 2H 2 O 

On exposure to air, sulfurous acid may be oxidized to sulfuric acid 
to some extent. 

2H 2 SO 3 + O 2 - 2H 2 SO 4 

In a similar way, sulfites may be oxidized to sulfates. 

How sulfurous acid, sulfites, and bisulfites are used. The tend- 
ency of one molecule of sulfurous acid to unite with one atom of 
oxygen, forming sulfuric acid, makes it a valuable reducing agent. 
This property of sulfurous acid enables it to be used in the bleach- 
ing of wood pulp, straw, silk, feathers, dried fruits, flour, molasses, 
and canned corn all of which would be partially destroyed by 
other bleaching agents. 

Fruits that are to be dyed are frequently bleached with sulfur 
dioxide (as sulfurous acicl) . For example, in making maraschino 
cherries, the cherries are first bleached and then dyed any desired 
color with edible dye. If sulfurous acid is used either in bleaching 
or in preserving a food that is sold in interstate commerce, this fact 
must be clearly stated on the label of the package in which it is sold. 

Chemical Construction Corporation 

A sulfur combustion furnace 
for the commercial production 
of sulfur dioxide. Melted sulfur 
is pumped into the furnace 
and burned to form the gas. 


Substances of vegetable origin are not easily destroyed by sulfurous 
acid, and therefore their coloring matter may be reduced to colorless 
compounds by this weak acid. However, the bleaching action of 
sulfurous acid is not as lasting as that of chlorine or hydrogen per- 
oxide, because the coloring matter that has been reduced may, on 
exposure to air, become reoxidized. This explains in part why straw 
hats, bleached with sulfurous acid, may turn brown. 

Calcium bisulfite, Ca (HSO 3 ) 2 , containing dissolved SO 2 , is used in 
converting ground wood pulp into sulftte pulp by dissolving out the 
gluelike substances found mixed with the cellulose of the wood pulp 
(see page 598) . Sulfite pulp is used in making paper. 

Sodium sulftte mixed with a solid acid, such as oxalic (white crys- 
tals) , makes up the dry powder used in bleaching straw hats. The 
action of this mixture of acid and sulfite illustrates the part that ions 
play in chemical reactions. While the dry solid acid and sulfite are 
in contact, no action occurs. As soon as water is added, however, the 
substances go into solution, dissociate, and the ions formed react, 
producing sulfurous acid, which bleaches. Na 2 SO 3 is also used as an 
antichlor (a substance used to remove any excess of chlorine after 
bleaching) and as an oxygen-removal agent in treating boiler-feed 

How to test for a sulfite. The characteristic od6r of the sulfur 
dioxide liberated when an acid is added to a sulfite constitutes a 
test for any sulfite. 

CaSO 3 + 2HC1 -> CaCl 2 + H 2 O + SO 2 T 

Preparation and properties of sulfur trioxide. When sulfur di- 
oxide is mixed with air and the mixture is passed over a catalyst 
such as heated platinum, sulfur trioxide is formed. 

2SO 2 + O 2 -> 2SO 8 

This constitutes the key reaction in the most important commercial 
method of manufacturing sulfuric acid, the contact process, so named 
because contact of the gases with this platinum catalyst plays so im- 
portant a part. 

Sulfur trioxide is a white solid, but it melts at about room temper- 
ature and boils at 46 C. The liquid fumes in moist air and reacts 
vigorously with water making a hissing sound, liberating much heat 
and forming sulfuric acid. Thus, sulfur trioxide is the anhydride of 
sulfuric acid. 

SO 3 + H 2 O - H 2 SO 4 


,"" for '"' 

! : : ; Mef 

Monsanto Chemical Company and Davison Chemical Corporation 

A contact process installation. Trace the process in this photograph and that on 
page 311. 

How sulfuric acid is made by the contact process. The manu- 
facture of sulfuric acid by the contact process may be conveniently 
divided into three stages: (1) burning sulfur to sulfur dioxide in 
a combustion furnace, (2) changing this sulfur dioxide to sulfur tri- 
oxide by passing sulfur dioxide and air over a catalyst in a converter, 
(3) changing the sulfur trioxide to sulfuric acid by passing the sul- 
fur trioxide into absorption towers through which sulfuric acid is 
circulated. The chemical reactions that take place are: 

1) S + O 2 -> SO 2 

2) 2S0 2 4- 2 -> 2S0 3 

3) SO 3 4- H 2 O [+ H ? SO4! - i2|H 2 SO 4 

1) The sulfur combustion furnace. In the combustion furnace, 
sulfur is brought into intimate contact with thoroughly dried air. 
A blower supplies this air at sufficient pressure to force it through 
the furnace and to force the sulfur dioxide from the furnace through 
the rest of the system. The heat produced by the combustion of the 
sulfur is recovered by means of a waste-heat boiler. Part of this heat 
is used in some plants to melt the sulfur that is sprayed into the 



2) The converter. The converter is simply a large chamber con- 
taining many perforated shelves covered with a catalyst in whose 
presence the union of warm sulfur dioxide and oxygen takes place, 
forming sulfur trioxide. Noteworthy developments in the contact 
process have been the introduction of vanadium pentoxide and 
platinized silica gel catalysts as substitutes for the more expensive 
platinum mass. The vanadium catalyst has the further advantage of 
being insensitive to poisoning by impurities, such as arsenic, anti- 
mony, selenium, and chlorine, that may be present in the sulfur. 

Since, technically, the greatest cost in processes that depend upon 
a catalyst is keeping the catalyst entirely free from minute traces of 
poisons, a catalyst less affected by such impurities is preferred. The 
introduction and use of the vanadium catalyst was soon followed by 
a platinum catalyst that is immune to arsenic poisoning and may 
be operated at a lower temperature. 

The change of sulfur dioxide to sulfur trioxide, which occurs on 
the surface of the catalyst, is an exothermic one, that is, heat is given 
off. Careful temperature control is necessary to insure highest yields. 
Below 450 C., some SO 2 escapes oxidation to SO 3 . Above 450 C., 
some SO 3 is decomposed to SO 2 . At a temperature of 450C. the 
reversible reaction 

2SO 2 + O 2 *= 2SO 3 

goes to the right almost completely. Only about three percent of the 
total amount of SO 3 theoretically possible is changed back to SO 2 . 

3) The absorption towers. From the converter, the sultur trioxide 
passes through a cooler, and then into a series of three or four ab- 
sorption towers. Through the first two of these towers, filled with 
quartz pebbles or acid-resisting packing rings, concentrated sulfuric 
acid is circulated. This concentrated sulfuric acid absorbs the sulfur 
trioxide, forming sulfuric acid of more than 98-percent concentra- 
tion. SO 3 vapor is absorbed more easily by concentrated H 2 SO 4 than 
by water. The third tower is a coke-filled filter tower, which absorbs 
any acid vapor that might otherwise escape into the air. 

A contact plant is practically automatic. Control of the process is 
maintained from one central room containing all the recording and 
controlling instruments. 

How the contact process compares with the lead-chamber process 
of preparing sulfuric acid. The older commercial process for making 
sulfuric acid is the lead-chamber process. It was introduced in Bir- 
mingham, England, in 1746 by Dr. Roebuck, a physician. Its use 



marked the beginning of chemical manufacture on a large scale, 
for the Industrial Revolution was just then getting under way. 

In recent years, the lead-chamber process has given way gradually 
to the contact process. Because of the large volumes of gases that 
must react (without a metallic catalyst) , large reaction chambers 
are needed. On account of the action of dilute su If uric acid on iron, 
it is necessary to line these large rooms with lead sheets, two inches 
thick a second item of expense. Some of the chemical reactions 
that occur in the chamber process are still not thoroughly under- 

Sulfur dioxide, steam, and the oxygen of the air are converted into 
dilute sulfuric acid by means of the oxides of nitrogen, which act as 
catalytic agents. The change of NO into NO 2 , which takes place in 
air, and the subsequent reversal of this reaction are indirect means 
of getting an extra atom of oxygen to combine with sulfur dioxide. 
At the end of the process a corrosive, sour drizzle of dilute sulfuric 
acid falls on the floor of the lead chambers. 

Another disadvantage of the lead-chamber process is that only 
dilute sulfuric acid, often impure and of not more than 78-percent 
concentration, can be made. If concentrated acid is required, this 
dilute acid must be concentrated by evaporation with heat. 

The chamber process can compete with the contact process only 
if acid of not greater than 78-percent concentration is required, as 
in the manufacture of phosphate fertilizers and in the pickling of 
steel. Sulfuric acid of less than 78-percent concentration must be 
shipped in glass or in lead-lined steel containers, while acid of 

In these towers the final steps of the contact process take place. 

Monsanto Chemical Company and Garfield Chemical and Manufacturing Corporation 



Fig. 70. Diluting concentrated 
sulfuric acid. Why is such care 
not necessary when diluting hy- 
drochloric acid? 



H 2 

H 2 S0 4 

greater than 78-percent concentration may be shipped in steel drums 
or tank cars. Why? 

Physical properties of sulfuric acid. Sulfuric acid, or oil of vitriol, 
is a water solution of hydrogen sulfate (a liquid) . Concentrated sul- 
furic acid contains 98 percent acid, has a specific gravity of 1.84, and 
boils at 338C. It mixes with water in all proportions, liberating 
much heat. For this reason, great care should be used in diluting 
concentrated H,SO 4 . The heavier sulfuric acid should be added to 
cold water slowly, and the mixture constantly stirred as shown in 
Fig. 70. Always add the acid to the water. Oleum, or fuming sulfuric 
acid, contains dissolved sulfur trioxide, and usually its formula is 
written H,SO 4 SO 3 . 

Chemical properties of sulfuric acid. A water solution of sulfuric 
acid is a very strong acid, because it dissociates to a high degree. The 
concentrated acid, because it does not dissociate to any appreciable 
extent, is weaker than the dilute acid. 

Because of its strong attraction for water, sulfuric acid is an excel- 
lent drying, or dehydrating, agent. Its attraction for water is so 
intense that sulfuric acid will remove the hydrogen and oxygen pres- 
ent in carbohydrates in which these elements occur in the same pro- 
portion as in water. Thus, when concentrated sulfuric acid is poured 

over sugar, C 12 H, 2 O n , it removes the H, 2 O U as 1 1 molecules of steam, 
and a black mass of pure carbon remains. 

- 11H 2 O + 12C 

Likewise, wood, which is composed of cellulose, C (1 H 10 O 5 , chars 
when clipped in concentrated sulfuric acid. Cotton, wool, and other 
substances react similarly. The dehydrating action of sulfuric acid 
accounts also for the severe burns it may produce on the skin. 

At 450C. and atmospheric pressure, sulfuric acid decomposes 
completely into water and sulfur trioxide. This chemical reaction 
is reversible. 

H 2 SO 4 ^ SO 3 + H 2 O 

It has been pointed out that nitric: acid is an oxidizing agent be- 
cause, when added to a metal, it forms water instead of hydrogen, 
and liberates a gaseous oxide. Warm concentrated sulfuric acid, 
which behaves in a manner similar to nitric acid, is also an excellent 
oxidizing agent. 

Cu + 2H 2 S0 4 -> CuS0 4 + 2H 2 O + SO 2 1 

This reaction between concentrated sulfuric acid and copper illus- 
trates another common laboratory method for preparing sulfur di- 
oxide. In contact with certain metals, including zinc, iron, and 
magnesium, dilute sulfuric acid liberates hydrogen as shown in the 
following equation, 

Mg + H 2 SO 4 - MgSO 4 +H 2 1 

Iron and steel (pickling) 

Rayon and film 

Other metallurgical 



Sulfuric acid aids agriculture. In a recent year, the United States 
produced about 50 percent of the sulfuric acid consumed by the 
entire world. The domestic consumption amounted to slightly more 
than 12 million tons. While it is not likely that most farmers believe 
themselves in any large measure dependent on sulfuric acid, it is a 
fact that the fertilizers used very widely in agriculture account for 
almost one-fourth of the sulfuric acid consumed each year. 

Phosphorus occurs in nature in a fairly plentiful mineral called 
tricalcium phosphate, or rock phosphate. Unfortunately, rock phos- 
phate is insoluble in water and therefore cannot be utilized by 
plants. However, when treated with H 2 SO 4 , rock phosphate changes 
into a soluble compound monocaldum phosphate (the active ingre- 
dient of superphosphate see page 487) . The phosphorus is then 
available as plant food. 

In cases where the soil is acid, however, pulverized raw rock phos- 
phate is applied directly. The acids of the soil attack the phosphate 
slowly and make part of it available as plant food. 

Sulfuric acid in production of petroleum products. The petroleum 
industries consume about one-eighth of all the sulfuric acid produced 
in the United States. Crude petroleum contains a large number of 
carbon compounds that are dark in color or that become dark from 
exposure to air. These compounds are removed from crude oil by 
treatment with sulfuric acid. In this way most of the waxy and 
gummy materials that clog burners and carburetors are removed. 

Other uses of sulfuric acid. Another important use of sulfuric acid 
is in the cleaning, or pickling, of sheet steel before covering the steel 
with a layer of zinc in galvanizing, or tin in tin-plating. The iron 
sulfate that is formed as a byproduct is crystallized and is used in 
making inks. The common name for crystallized iron sulfate is 
copperas, once known as green vitriol. 

Fe + H 2 SO 4 + 7H 2 O - H 2 1 + FcSO 4 7H 2 O (copperas) 

The manufacture of explosives is dependent upon large quanti- 
ties of concentrated sulfuric acid. It is used as a dehydrating agent 
in the process of nitrating the many substances from which explo- 
sives are made. Sulfuric acid is used in the production of textiles and 
rayon and cellulose film. Many acids, including hydrochloric and 
nitric, are made with the aid of sulfuric acid. Others of the hundreds 
of processes in which sulfuric acid is used include the making of sul- 
fates, paper, rayon, leather, celluloid, dyes, dn$$p, and paints. (Many 
of these processes will be described later.) Sutewj&acid is used also 
as the electrolyte in lead storage batteries. ^%r 


Normal, acid, and basic salts. Zinc sulfate, ZnSO 4 , is formed by the 
reaction of zinc and sulfuric acid. Here both of the hydrogen ions 
of the acid have been replaced by the metal, and the resulting salt 
is a normal salt, that is, a salt in which all the hydrogen of the acid 
has been replaced by a metal. In addition to forming normal salts, 
acids such as H 2 SO 4 , H 2 CO S , and H 2 SO 3 , which contain two hydro- 
gen atoms, can also form acid salts, and are called dibasic acids. Boric 
acid, H 3 BO 3 , which contains three replaceable hydrogen atoms, is a 
tribasic acid. 

In the preparation of hydrochloric acid by the action of sulfuric 
acid on sodium chloride, sodium bisulfate, NaHSO 4 (also called so- 
dium acid sulfate) , is obtained as one of the byproducts. 

NaCl + H 2 SO 4 -> NaHSO 4 + HC1 

Here only one of the hydrogen ions has been replaced, and the other 
remains in the salt, which is called an acid salt. Thus it can be seen 
that either one or both of the hydrogen ions of sulfuric acid may be 
replaced by a metal. 

Acids such as HC1 and HNO a contain only one hydrogen atom and 
are called monobasic acids. A salt containing one or more hydroxide 
groups is called a basic salt. Such a salt is basic lead carbonate 
Pb 8 (OH) 2 (CO,) a (see page 449) . 

Glauber salt and epsom salt. One of the most famous chemists of 
the seventeenth century was a Bavarian, Johann Glauber. When he 
was 21 he was attacked by a fever and was advised to drink the water 
of a certain well. After recovering from his illness he analyzed wa- 
ter from this well, and extracted from it crystals of a salt which he 
called sal mirabile, the wonderful salt. He recommended it as a 
"splendid medicine for internal and external use." This salt, 
Na 2 SO 4 . 10H 2 O, known as glauber salt has been used as a laxative 
for more than 300 years. Anhydrous Na 2 SO 4 , also known as salt cake, 
is used in the manufacture of paper and glass. 

Soon after the introduction of glauber salt, an English physician, 
Nehemiah Grew, who discovered that plants have sex, extracted 
a salt from some springs in the village of Epsom, near London, 
England. He wrote a book on the medicinal value of this salt, 
MgSO 4 7H 2 O, and thus epsom salt, a white, soluble, crystalline 
compound became a rival laxative. 

Many proprietary medicines, mineral-water crystals, and mineral 
waters sold today at ridiculously high prices are laxatives containing 
one or both of these salts. Such salt mixtures and solutions are sold 
to be used, often dangerously, as "the modern way to slenderize." 


Reducing preparations of this kind should not be taken without the 
advice of a competent physician. 

How to test for a sulfate ion. Most sulfates are soluble in water. 
However, the sulfates of barium, lead, and strontium are insoluble. 
When barium chloride is added to a sulfate, a white precipitate, 
barium sulfate, is formed. This white precipitate is distinguished 
from other barium salts by the fact that barium sulfate is insoluble 
in hydrochloric acid. 

Na 2 SO 4 + BaCl 2 - BaSO 4 J + 2NaCl 
This reaction goes to completion. Why? 


Aaron, Harold. Good Health and Bad Medicine. Consumers 
Union of United States, Inc., New York, 1940. Carefully pre- 
pared and well-written materials on self-medication. 

Goldblatt, L. A., Ed. Collateral Readings in Inorganic Chem- 
istry. D. Appleton-Century Co., New York, 2nd series, 1942. 
Article 15 contains an excellent description of a "Contact 
Sulfuric Acid Plant." 

Leicester, H. M., and Klickstein, H. S. Source Book in Chem- 
istry. McGraw-Hill Book Company, New York, 1952. Pages 
11-16 contain Georg Agricola's description of the manufacture 
of vitriol, from De re Metallica. 


Group A 

1. Copy and complete the following statements. Do not 
write in this book. SO 2 is prepared for industrial use by burn- 
ing ... or roasting .... It occurs in nature in small amounts 
near .... The most important use of SO 2 is in the manufacture 
of .... Its most characteristic physical property is .... 

2. (a) What two methods are used in preparing SO 2 in 
the laboratory? (b) Which method involves reduction? (c) Ex- 

3. (a) Make a labeled diagram showing how SO 2 is pre- 
pared from Na 2 SO 3 . (b) Write the equation for the reaction 
that occurs. 

4. Compare the water solubility of SO a ,*faH 3 , HC1, and O 2 . 


5. Write three equations illustrating three chemical prop- 
erties of SO 2 . 

6. Describe three uses of SO 

2 . 

7. For what purpose is SO 2 used in the household? 

8. (a) Is SO 2 formed first by the action of H 2 SO 4 on 
Na 2 SO 3 ? (b) Explain. 

9. (a) Write the reversible reaction between SO 2 and 
H 2 O. (b) How can the reaction be made to go to the right? 
(c) to the left? 

10. (a) Why is H 2 SO 3 a weak acid? (b) What is a sulfite? 
1 1 . Write an equation showing the formation of Na 2 SO 3 . 

12. (a) By what chemical process does H 2 SO 3 bleach? 
(b) Into what does H 2 SO S change as it bleaches? (c) Compare 
the way in which H 2 SO 3 bleaches with the way in which C1 2 

13. Write H 2 SO 3 , H 2 O 2 , and C1 2 in a column. Opposite each 
write the names of the substances that it is best suited to 

14- (a) Why is bleaching with H 2 SO 3 less permanent than 
bleaching with other chemicals? (b) Illustrate. 

15. Distinguish between a sulfite, a sulfate, and a sulfide. 

16. How much S would be needed to prepare 500 tons of 
98 percent H 2 SO 4 ? 

17. Cu is added to concentrated H 2 SO 4 and 448 ml. of SO 2 
are liberated. How much H 2 SO 4 is decomposed? 

18. Calculate the percentage of S in each of its two oxides. 

19. What are three properties of SO 3 ? 

20. Write the equation for the burning (roasting) of FeS 2 . 

21. What are the stages in the contact process of making 
H 2 S0 4 ? 

22. Write equations that show the three chief chemical 
changes that take place in the preparation of H 2 SO 4 by the 
contact process. 

23. (a) Why must the catalyst used in the contact process 
be chosen with special care? (b) What advantage has V 2 O 5 
over the old plantinum catalyst used in the manufacture of 
H 2 S0 4 ? 


24. Describe briefly the chemical change that is the basis 
of the chamber process for making H 2 SO 4 . 

25. (a) What are the disadvantages of manufacturing 
H 2 SO 4 by the chamber process? (b) This process competes 
favorably with the contact process in producing H 2 SO 4 for 
what industry? (c) Why? 

". i " " " ; 

26. What are the physical properties of oil of vitriol? 

27. In mixing concentrated H 2 SO 4 and H 2 O, why should 
the acid always be added to the H 2 O rather than the H 2 O to 
the acid? 

28. Devise an experiment to show the dehydrating action of 
H 2 S0 4 . 

29. By its action on Cu, show that concentrated H 2 SO 4 is 
an oxidizing agent. 

30. (a) Make a list of the uses of H 2 SO 4 . (b) From the 
viewpoint of the amount consumed, what is its chief use? 
(c) Explain fully, using an equation to show the reaction in- 

31. Why does the refining of petroleum require immense 
quantities of H 2 SO 4 ? 

32. Name three experiments in which you used H 2 SO 4 in 
making laboratory preparations. 

33. What is sulfite pulp? 

34. What is the difference between a normal, a basic, and an 
acid salt? 

35. (a) Name three sulfates and (b) state an important use 
of each. 

Group B 

36. Why is SO 3 not added directly to water in making 
H 2 SO 4 in the last stage of the contact process? 

37. Make a diagram of the apparatus you would use in pre- 
paring SO 2 in the laboratory by the action of H 2 SO 4 on Cu. 

38. Explain: (1) Black rings often form on wooden shelves 
that hold bottles of H 2 SO 4 . (2) A full battle of concentrated 
H 2 SO 4 overflows when exposed to air. (3) An open bottle of 
concentrated H 2 SO 4 is sometimes placed in a desiccator. 
(4) Frightful burns result from getting concentrated H 2 SO 4 on 
the hands. * 


39. Compare the three strong acids with respect to their 
action on Cu and Zn. Consider both dilute and concentrated 

40. A load of scrap iron weighing 2715 pounds and con- 
taining 95.5% iron is added to a large vat of dilute sulfuric 
acid. What is the maximum amount of ferrous sulfate that can 
be obtained? 


1. If you happen to have epsom salt in your medicine chest, 
bring it to school, read the label to your class and make a test 
for a sulfate with some barium chloride solution your teacher 
will supply. 

2. Write a two- or three-page report to justify the statement 
that H 2 SO 4 is "king of chemicals," and the "basis of the 
machine age." Prepare a debate or a class discussion based on 
these statements. 




. . . It (the light from my lamp) 
was hurled out of the sun 200,000,- 
ooo years ago, and was captured by 
the leaves of the Carboniferous tree- 
fern forests, fell with the falling 
plant, was buried, fossilized, dug up 
and resurrected. It is the same light. 
And, in my little fig tree as in the 
ancient ferns, it is the same unchang- 
ing green stuff from age to age. 
Donald Culross Peattie in Flowering 
Earth, 1939. 

The fascinating chemistry of carbon. Carbon is the eleventh most 
abundant element of the earth by weight. In the form ot coal and 
graphite, it helps to drive as well as lubricate the wheels of our 
machine age. As petroleum, it propels our locomotives, ships, auto- 
mobiles, and airplanes. In the form of diamond and "Carborundum" 
(synthesized directly from free carbon) , carbon is used in making 
tools that drill our hardest rocks and grind and polish our machines, 
tools, and household utensils. 

In combined carbon we find one of the most fascinating stories 
in the whole romance of chemistry. Each year carbon's 650,000 
known compounds are being increased in number by thousands of 
compounds synthesized in our industrial and university laborato- 
ries. Synthetic chemistry reaches its greatest development in organic 
chemistry, the chemistry of the compounds ot carbon. 

Woehler (vu-ler) , in 1828, synthesized urea, the first organic com- 
pound made outside the living body. By this synthesis the old idea 
that organic compounds could be formed only in living matter was 
destroyed, and a new era in chemistry was ushered in, witnessing 
"the great tragedy of science, the slaying of a beautiful hypothesis 
by an ugly fact." 




One hundred and twenty-five years after Woehler's achievement, 
sugar (C^H^O^) was synthesized from its elements. What a century 
and a quarter of research between Woehler's urea and synthetic 
cane sugar! Berthollet went to the red ant and learned the secret of 
preparing formic acid, the liquid responsible for the sting of this in- 
sect. William Perkin, washing bottles in a London laboratory, mixed 
at random the contents of two bottles and discovered a method of 
synthesizing mauve the first of a long series of organic dyes that 
rival the colors of nature. And this is only the beginning. The mind 
fairly reels at the accomplishments and possibilities of this branch of 
chemistry. More of this entrancing story is told in later chapters. 

What is the valence of carbon? Carbon has an atomic weight of 
12. The nucleus of the carbon atom contains six protons; hence, the 
outer ring contains only four electrons. Therefore, the carbon atom 
must either lose four electrons or gain four electrons to complete its 
outer ring. The valence of carbon then is either plus four or minus 
four, depending upon the substance with which it reacts. That is, 
it may act as either a metal or a nonmetal. Carbon almost always 
forms firm covalent bonds involving four pairs of shared electrons. 

The properties of carbon. Carbon occurs in three allotropic forms: 
two of them, diamond and graphite, are crystalline; the third is 
amorphous, or noncrystalline, in form. All forms of carbon combine 
with oxygen, forming either carbon monoxide, CO, or carbon di- 
oxide, CO 2 . The different forms of carbon vary widely in the ease 
with which they may be oxidized. Coal burns easily particularly 
when powdered. On the other hand, diamond combines with oxygen 
only at extremely high temperatures. 

C + O2 > CO2 (carbon dioxide) 
2C + O 2 > 2CO (carbon monoxide) 

Its attraction for oxygen makes carbon a good reducing agent. The 
reducing action of carbon may be shown by heating a mixture of 


Fig. 71. Blowpipe analysis 
of an ore (left) using the re- 
ducing power of charcoal, 
(right) The carbon atom. 


carbon and lead oxide on a charcoal block, using a blowpipe. A lit- 
tle ball of pure lustrous lead is formed and CO 2 is given off. 

2C 4- Pb 3 O 4 -> 3Pb + 2CO 2 T 

Millions of tons of coke (carbon) are used industrially every year 
in reducing iron oxide to commercial forms of iron. 

Fc 2 O 3 + 3C -> 2Fc + SCO | 

This reaction shows carbon acting as a metal, a lender of electrons. 
Carbon may also behave as a nonmetal, a borrower of electrons. At 
the very high temperatures reached in an electric furnace, carbon 
unites with metals such as calcium and wolfram and with less electro- 
negative elements such as silicon, forming carbides. 

CaO +3C-> CaC 2 + CO t 

calcium oxide calcium carbide 

SiO 2 +3C- SiC +2COT 

silicon dioxide silicon carbide 

Carbon unites directly with hydrogen also at the high temperatures 
of an electric arc, forming acetylene gas, C..H.,. 

2C + H 2 -> C 2 H 2 T 

All forms of carbon are insoluble in water, acids, and bases. 
Amorphous carbon, however, is soluble in some molten metals in- 
cluding iron. 

Physically, the allotropic forms of carbon resemble one another 
in some respects. They are all odorless, tasteless, and insoluble in 
water. On heating, they do not melt but vaporize. 

The hardest substance known. In the eighteenth century, diamond 
was shown by Lavoisier to be almost pure carbon. This French chem- 
ist burned a diamond in the presence of a distinguished audience by 
concentrating the sun's heat on it with a lens. He showed that the 
gas which was formed was pure carbon dioxide. This startling dis- 
covery strengthened Lavoisier's theory of burning. Why? 

Diamond, chief among precious stones, is found mainly in South 
Africa, Brazil, and Borneo. It occurs in various sizes and states of 
purity, buried deep in the ground frequently in the pipes of extinct 
volcanoes, or in the loose sands along certain streams. The brilliance 
of diamond, caused by its power of refracting light, is brought out 
by removing any surface impurities or imperfectiqns and by cutting 
facets on its surface. Small amounts of metallic oxides in diamonds 



Shell Oil Company Baunigold Bros., Inc. 

This asphalt core, neatly cut from a Before use in jewelry, a diamond is 
roadbed by a diamond drill, will be polished against a rotating iron disc 
taken to a laboratory for study. coated with olive oil and diamond dust. 

give them various colors, the rarest of which are red, green, and blue. 
The density of diamond is 3.5. 

Diamond is weighed in carats, a unit of weight based on the weight 
of the bean of the carob tree. One carat is equal to one-fifth of a 
gram. The largest diamond ever mined was the flawless Cullinan 
found in South Africa in 1905. It weighed 3026 carats, or 21.35 
ounces, and was about the size of a man's fist. 

From the first little diamond found in a pebble picked up by a 
child in 1867, the diamond industry of South Africa has grown to 
world leadership. In a recent year, more than three million carats of 
diamonds were mined. About half of this amount is used in jewelry. 

Though brittle, diamond is the hardest substance known, and for 
this reason is used extensively in making tools for drilling rock, 
grinding, sawing, engraving, and polishing/ Impure black diamonds, 
called carbonado and bort, are used in drills.) Diamonds are also used 
as bearings in watches and delicate instruments, and in wire-drawing 
dies made by drilling tapered holes through diamond crystals. 

Can diamonds be made synthetically? Scientists believe that 
diamonds were formed under great heat and pressure by the crystal- 
lization of carbon dissolved in molten iron or other substances. To 
test this theory of the formation of diamonds, Henri Moissan (mwa'- 
sa'N') , in 1893, subjected some pure carbon to extremely high tem- 
peratures and pressures. He dissolved carbon in molten iron, and 
quickly plunged the molten mass into cold water. The sudden 
contraction of the iron exerted tremendous pressure on the dissolved 


carbon, causing it to crystallize. The iron was dissolved with acids, 
and, aided by a microscope, Moissan reported seeing synthetic dia- 
monds formed of tiny cubical crystals. Other attempts have been 
made to synthesize diamonds, but thus far these attempts have not 
proved a threat to the diamond industry. 

Black, soft, and greasy. A second allotropic form of carbon, 
graphite, is lustrous, black, soft, and oily to the touch. The great 
difference in hardness between the allotropic forms of carbon is 
caused by the different arrangement of the atoms. The atoms of the 
diamond molecule are very closely interlaced and bound firmly to- 
gether. Graphite, on the other hand, is made up of piles of flat lay- 
ers of atoms firmly bound to one another, but far apart and only 
feebly hanging on to the next layer. For this reason graphite is soft 
and slippery. The flat planes slide over one another like cards in a 
pack. A few electrons between the planes are loosely held, making 
graphite an excellent conductor, whereas diamond is a nonconductor. 

Natural graphite, which is found chiefly in Ceylon, Siberia, 
and Canada, contains many impurities. It is believed to have been 
formed, like diamond, under high temperature and pressure. 

Electric furnaces and uses of synthetic graphite. In 1896 Edward 
Acheson, ot Pennsylvania, patented a process for making synthetic 
graphite,, today a widely used article of commerce. It is superior to 
natural graphite, because it is free from grit and contains almost no 
impurities. Acheson was at different times a blast-furnace helper, oil 
gager, raih*oad ticket agent, store clerk, miner, bookkeeper, and en- 
gineer. He also worked for Edison and was sent abroad to the Inter- 
national Exposition in Paris to exhibit the Edison electric lamp 
which he had helped to improve with a graphite filament. His 
method of making synthetic graphite consists of subjecting anthra- 
cite coal or molded carbon to the high temperature of an electric 
furnace, keeping air out by covering the coal with sand. 

An electric furnace, as you know, transforms electric energy into 
heat energy. The resistance to an electric current offered by loose 
carbon, in the center ot the resistance type of electric furnace used 
in making graphite, produces high temperatures. Other electric fur- 
naces utilize the heat produced by the passage of electricity through 

carbon atoms 

Fig. 72. Diagram of a graphite 
particle showing the layers of 


Alloy steels are produced 
in this electric furnace. The 
electrodes of the furnace 
are made from graphite. 

Electro Metallurgical L'o mint try 

high-resistance wire, such as an alloy of nickel and chromium. An- 
other type of electric furnace is the arc type. Here the passage of 
electricity across an air gap produces a high temperature as a result 
of the resistance of the air. Arc-type electric furnaces are used in the 
synthesis of nitric oxide, NO, during the preparation of nitric acid 
by nitrogen-fixation. 

Like other forms of carbon, graphite does not melt and has an ex- 
tremely high point of vaporization. For this reason it is used in mak- 
ing graphite crucibles in which steel and other alloys of high melting 
point are prepared. Because of its smoothness, oilincss, and high tem- 
perature of fusion, graphite is mixed with oils used in lubricating 
heavy, swiftly moving parts of machinery, in which even heavy oil 
would evaporate or burn off. 

Graphite electrodes are used in electric furnaces, arc lamps, and 
in general, for conducting electricity at very high temperatures. 
Powdered graphite is used in dry cells, and in certain oil-retaining 
bearings produced by powder metallurgy (see page 467) . 

Graphite is used also in lead pencils. The "lead" in a pencil is 
graphite mixed with clay to vary the hardness of the lead. Graphite 
is also a constituent of stove polishes. Most synthetic graphite is pro- 
duced at Niagara Falls, where hydroelectric power is available at 
low cost. 


C0dl v : | y Fig 73 Cross section of an 

electric furnace used in the 
commercial production of 
synthetic graphite. 

graphite electrodes 



"Formless," yet of many forms. The noncrystalline allotropic form 
of carbon is amorphous carbon. The purest form of amorphous car- 
bon can be prepared easily by heating pure table sugar. The hydro- 
gen and oxygen are driven off as water; pure amorphous carbon is 

Ci2H 22 Oii -> 11H 2 O + 12C 

Among the most common forms of amorphous carbon are (a) lamp- 
black and carbon black, (b) gas carbon, (c) coke, (d) charcoal, and 
(e) animal charcoal, or boneblack. 

Valuable soot. Lampblack was known to the Chinese and Egyp- 
tians who used it in making ink. It is amorphous carbon dust formed 
by the incomplete burning of compounds rich in carbon. It is man- 
ufactured commercially by burning heavy liquid hydrocarbons (com- 
pounds of hydrogen and carbon) in an insufficient supply of oxy- 
gen. v The escaping particles of carbon dust, or soot, from the small 
flames are collected on a revolving metal plate that is kept at a low 
temperature by flowing water. The smokiness of a kerosene lamp 
burning with an insufficient air supply or with the oil supplied too 
rapidly illustrates the way in which lampblack is formed. 

Carbon black is made in the same way as lampblack except that 
natural gas is used instead of liquid hydrocarbons. It* is finer than 
lampblack, which it is gradually replacing. Carbon black is such an 
important item of commerce that in a recent year its production ac- 
counted for more than ten percent of all the natural gas produced in 
the United States. Most of the carbon black manufactured in this 
country is added to rubber tires to improve their resistance to tear 
and abrasion. It is also used as a basic raw material in making print- 
er's ink, typewriter ribbons, carbon paper, phonograph records, and 
black paints. 

Almost graphite. Gas carbon is amorphous carbon that collects on 
the walls of retorts during the manufacture of illuminating gas from 
coal. In structure, gas carbon is semicrystalline and resembles graph- 
ite. It is used in the rods of arc lamps and other electrodes. 

Coke from the destructive distillation of coaL Heating a complex 
carbon compound in an oven from which air is* excluded, and con- 
densing the vapors formed is called destructive distillation. When 
soft coal is so treated the most important solid product is coke, a 
steel-gray, hard, brittle substance used chiefly as a reducing agent 
in the extraction of iron, and as a fuel. The coal is heated to 2000 F. 
for 16 hours, and then the white-hot coke is pushed out from the 
oven and quickly quenched in cold water. 


In the byproduct coke oven many valuable chemicals are recov- 
ered. As the gases pass through oil-absorption tanks, a light oil sepa- 
rates, from which benzene, toluene, xylene, and naphthalene are dis- 
tilled. From the coal tar which is formed anthracene, phenol, and 
pitch are also separated by fractional distillation. Ammonia gas and 
coal gas are also obtained (see pages 362-363 for a detailed account 
of this process) . Less than ten percent of our coke is still made in 
beehive ovens from which none of these byproducts are recovered. 

Charcoal-broiled. If wood is subjected to destructive distillation, 
charcoal is one of the products. Charcoal retains the fibrous struc- 
ture of wood. 

Powdered charcoal is frequently put together under tremendous 
pressure, forming compact, pillow-shaped blocks called charcoal 
briquets. These burn with no smoke and leave very little ash. They 
are used in camping or picnicking, on pleasure boats, and in found- 

Charcoal-broiled meats, especially steaks and chops, are particu- 
larly delicious. Meats broiled over charcoal do not pick up odors or 
flavors from the burning fuel. 

Activated charcoal. This is a specially-prepared charcoal used for 
removing or adsorbing gases. Adsorption is caused by the collection 
of thin layers of molecules of gases and other impurities on the sur- 
faces of the porous charcoal. It is not a chemical union but a physical 
attraction. Ill-smelling colored liquids, passed through such charcoal, 
are cleared of impurities responsible for the odor and color. Adsorp- 
tion is used also in recovering many industrial solvents and waste sub- 
stances, in air purification, and to concentrate a desired substance 
such as the drug streptomycin from its mold culture. 

Activated charcoal is used in gas masks to adsorb poison gases met 
in industry. "Nuchar" is the trademark for an activated carbon of 
vegetable origin containing about 90 percent pure carbon. 

A black decolorizer. The chief solid left after the destructive 
distillation of clean cattle bones is a black powder called boneblack. 
This boneblack, or animal charcoal, is a mixture containing only 
about ten percent amorphous carbon and 80 percent calcium phos- 
phate, the chief compound found in bones. It, too, has great adsorp- 
tive powers and is used mainly in the refining of crude sugar, from 
which it removes impurities that cause objectionable colors and 
odors. It is used also as a black pigment in paints. 

How coal was formed. Coal is found in all continents, including 
Antarctica, but the largest coal-producing areas are in North Amer- 
ica and Asia. Scientists now believe that many millions of years ago, 


during the Carboniferous period, great portions of the earth were 
covered with a dense vegetation more luxuriant than that found 
today even in tropical jungles. As the level of the earth sank during 
one of the many upheavals that occurred, vast portions of these 
jungles were flooded by the ocean and became swamps. Later these 
swamps were completely submerged and then gradually covered with 
mud, sand, or clay deposited as sediment by streams and rivers. 

Partial decomposition of the wood and other vegetable matter, 
aided by bacteria and fungi, changed the woody material first into 
peat. Minnesota, Florida, and several other states contain large peat 
deposits. Near Chester, Wisconsin, is a peat deposit that covers 
32,000 acres. Its average depth is six feet and it contains about 
40,000 tons of air-dried peat. Peat bogs also occur in Ireland and else- 
where. When dried and pressed, peat may be used as a fuel. It burns 
with a great deal of smoke and produces little heat. 

Scientists believe that the next step in the formation of coal was 
the changing of the peat into brown or black lignite^ which still re- 
tains the structure of the plants from which it was formed. Some- 
times branches or twigs are found in lignite in the form in which 
they grew. Lignite is found in some of the states west of the Missis- 
sippi, and, where no other form of coal is available, is used as a 
household fuel. In North Dakota, for example, lignite is being 
mined at the present time and utilized in making gas and briquet 
fuels, and for the production of electric power. 

Further decomposition and pressure, away from air, drove out 
more oxygen and hydrogen from the lignite, leaving hydrocarbons 
and some free carbon and forming bituminous, or soft, coal. Further 
heat and greater pressures changed the soft coal to jet black, lus- 
trous hard coal, or anthracite. 

Strong evidence supports this theory of coal formation. The fos- 
sils embedded in coal deposits include forms of both animal and 
vegetable life. Some of the animals appear to have been marine. 
Among the plants, mosses and ferns occur most commonly. 

The chief kinds of coal. The chief kinds of coal used in this 
country are soft, or bituminous, coal, and hard coal, or anthracite. 
Most of the anthracite mined in the United States comes from Penn- 
sylvania, yet this state, together with West Virginia, produces more 
than half the bituminous coal mined in this country. For every ton 
of anthracite mined, 11 tons of bituminous coal are dug. Anthra- 
cite contains practically no free carbon and produces less volatile 
material than bituminous coal (see Table 11) . It is cleaner to han- 
dle and burns with a short, pale blue flame. 


The use of anthracite is confined largely to the northeastern states, 
because these states have the advantage of being near the deposits 
and thus have short and low-cost transportation from the mines. In 
other parts of the country anthracite is relatively unfamiliar. 

The problem of smog. Smog, a mixture of smoke and fog, is 
formed largely by the incomplete combustion of coal, fuel oil, gas 
and rubbish. The products of this incomplete combustion are oxi- 
dized by the o/one in the air to create the smog. Smog may be aggra- 
vated by natural particles in the air, such as dust or pollen, or even 
by local weather conditions. 

People who live in cities where smog is a problem spend far more 
for soaps and other cleansers than might be necessary if better smoke 
control were practiced. It has been shown in several cities that smog 
may be eliminated. However, success in reducing smoke depends 
upon the cooperation of every consumer of fuel from the largest 
industry to the individual homeowner. 

Smoke control is particularly necessary in cities in which bitumi- 
nous coal is a major fuel. Bituminous coal, when burned in an in- 
adequate supply of oxygen, produces a very smoky flame. If the coal 
is properly fired (sec page 35!)) , smoke may be greatly reduced. 

British thermal units. A chief factor in determining the price of 
coal is its fuel value. This value is measured either in calories or in 
British thermal units, Btu. One Btu is the amount of heat required 
to raise the temperature of one pound of water one degree Fahren- 
heit. One Btu is equivalent to 252 calories. Burning one pound of 
average anthracite liberates about 12,700 Btu; burning the same 
weight of bituminous coal liberates about 13,100 Btu. The per- 
centage of ash and free sulfur is another factor important in deter- 
mining price. In general, anthracite contains a smaller percentage 
of these materials than does bituminous coal. 

A scene in Pittsburgh, Pa., before (left) and after a smog-control program was 

Allegheny Conference on Community Development 

TABLE 11. fl 




Variations in fixed || 

carbon, volatile matter, ^ 

and moisture on an ash ^ 

free basis of the several || 

ranks of coal produced ^ 

in the United States || 



Subbituminous C 

Subbituminouj B 

Subbituminous A 

High -Volatile C Bituminous 

High -Volatile B Bituminous 

High -Volatile A Bituminous 

Medium -Volatile Bituminous 

Low -Volatile Bituminous 



Meta- anthracite 




Coal, a major source of power. Coal has been used as fuel for 
centuries. Marco Polo, 700 years ago, reported the Chinese burning 
"black stone dug out of mountains where it runs in veins." But it was 
not widely used until the sixteenth century. By 1661 "a hellish and 
dismall cloud of Sea-Coale hung perpetually over London," where 
coal was used to make iron, glass, and other products. During the 
Industrial Revolution it became a vital source of energy. 

Our civilization uses energy at a staggering rate. In 1900, the total 
supply of energy from mineral fuels and water power in the United 
States was 7.4 quadrillion Btu; by 1950, the total had reached 36.2 
quadrillion Btu an increase of almost 400 percent! Each year our 
growing population and industries need added amounts of energy. 
It is expected that before the end of the century the annual produc- 
tion of energy will be more than double the present rate. 

Today, petroleum, coal, natural gas, and hydroelectric dams are 
our major sources of energy. In a recent year, petroleum supplied 40 
percent of our nation's total energy requirement; coal supplied 34 
percent; natural gas, 22 percent; and water power, 4 percent. In 
the years ahead, the use of nuclear energy will, in all probability, 
change these figures radically. 

Coal has many uses. Production of anthracite coal is now about 
40 million tons per year. About three-fourths of this, output goes to 
the retail market to be used as fuel for commercial and institutional 
buildings, as well as for heating homes. The remainder of the yearly 
tonnage is used by a variety of industrial consumers. 

Annual production of bituminous coal is now about 450 million 
tons per year. The major use of this coal is in generating electricity. 




About one-half of our nation's total electric power is produced by 
coal-fired steam plants. In all, the electric utilities consume consider- 
ably more than 100 million tons of coal each year, or about 25 per- 
cent of all bituminous coal marketed. The production of electric 
power is expanding rapidly and it is expected that there will be an 
increase in the amount of coal used for this purpose. 

The steel industry is the second most important customer for 
bituminous coal, consuming about 20 percent of all coal marketed. 
Most of this enormous quantity of coal is converted into coke by 
roasting in coke ovens. Coke is used to some extent as a fuel, but its 
chief use by far is as a reducing agent in the production of steel. 
Coke robs the iron ore, an oxide of iron, of its oxygen and leaves pig 
iron which is later converted to steel. The chemicals recovered in 
the coking process (see page 327) are the source of many plastics, 
dyes, drugs, medicines, and industrial chemicals. Today, coal chemi- 
cals are also produced directly from the "raw" coal without any at- 
tempt at burning it or changing it to coke. This use of coal as yet 
constitutes a very small market, but it may become increasingly im- 
portant in the future. Coal is also destructively distilled in the manu- 
facture of coal gas. 

All other types of manufacturing consume another 15 to 20 per- 
cent of the bituminous coal marketed annually. A similar amount 
goes to the retail market for use as a fuel in public buildings, com- 
mercial establishments, and homes. In addition to coal in its conven- 
tional lump form, the retail market each year uses over a million 
tons of coal in the form of fuel briquets. Fuel briquets are pressed 
cubes, cylinders, or ovoids of very fine bituminous coal (slack) or lig- 
nite held together by a burnable binder. 

The railroads, once the coal industry's best customer, now use 
considerably less than ten percent of the total annual tonnage as the 

Consumers Power Company 

The importance of coal in 
generating electricity is illus- 
trated by this coal storage 
pile at a modern power 
station. This plant burns 
about 675,000 tons of coal 
each year. 


coal-burning "iron horse" has been displaced by the more efficient 
diesel-electric locomotive. 

Buying coal for your home. The most important consideration is 
the amount of heat the coal will produce. Your coal dealer usually 
knows the Btu value of the coal. Such information may also be ob- 
tained from the United States Bureau of Mines, from your state bu- 
reau of mines, or from the company that mined the coal. 

If you know the number of Btu in a pound of coal, the product of 
this number multiplied by 2000 is the number of Btu per ton. The 
quotient of the number of Btu per ton divided by the cost per ton in 
dollars is the number of Btu per dollar. As far as heat value alone is 
concerned, the coal with the greatest number of Btu per dollar is the 
best buy. Expressed as a formula, this method of comparing coals 
may be written: 

Btu per Ib. X 2000 _ , 

- : rr; = Btu per dollar 

Cost per ton in dollars r 

However, as anyone who has ever carried out "ashes" knows, the ash 
content of coal is an important consideration also. One coal with a 
slightly higher Btu-per-dollar value than a second coal actually may 
not be as desirable as the second coal, if its ash content is appreci- 
ably higher. Moisture content and dust content are two other fac- 
tors that must be considered. Coal dealers may control the dustiness 
of coal either by sprinkling the coal with water or by spraying the 
coal with special oils. From the point of view of the consumer, which 
of these methods is preferable? Why? 

In buying coal, it is advisable to buy the size and grade recom- 
mended by the manufacturer of the equipment in which you intend 
to burn it. Buying the appropriate size and grade of coal will enable 
you to get more Btu out of the coal you buy. Most coal-burning 
equipment is designed to operate most efficiently with coal of cer- 
tain size and grade. 

Because of seasonal demands, coal is usually lower in cost during 
summer than in winter. Since coal does not deteriorate appreciably 
in dry storage, for example, in a coal bin, it is good policy to buy 
next winter's supply in the summertime. 

Changing energy relationships. The sources of energy upon which 
our civilization depends are, as we know, coal, petroleum, natural 
gas, and water power. Nuclear energy holds great promise, but its 
use as a major source of energy is still in the future. 

Petroleum and natural gas constitute but a tiny part of our na- 
tion's total fuel deposits. Taken together, they m^te up less than 


two percent of the known reserves of all fuels. Of the remaining 98 
percent, oil shale forms eight percent and coal forms 90 percent. 
More than 92 percent of the total coal reserve is bituminous and 
lignite; the remainder is anthracite. Although the experts disagree 
on the exact number of years involved, it is generally accepted that, 
at our present rate of use, our reserves of high-quality petroleum and 
natural gas may be consumed in less than half a century, while our 
coal reserves are adequate for several hundred years. Oil shale is an 
unexploited resource and we do not yet know the extent of its use- 

Because of coal's abundance, it seems likely that it will see in- 
creased use in the future as oil and gas reserves shrink. Because our 
society is completely dependent upon a plentiful and continuous sup- 
ply of energy, proper use and conservation of our fuel reserves is a 
problem which deserves the thoughtful attention of every citizen. 

The energy of running water, while limited in amount, is never- 
theless almost limitless in duration. For, in harnessing such energy, 
nothing is actually used up that will not occur again. The water cycle 
takes care of this problem for us. However, water power suffers from 
a serious drawback in that it is not constant. That is, during certain 
seasons of the year, more water is available than at others. Even 
dams and huge reservoirs have not yet succeeded in solving this prob- 
lem completely, but they have made tremendous strides in doing so, 
and more and more water power is being harnessed as the years go 
by. However, even if developed to the fullest extent, our water re- 
sources could not provide more than a small fraction of our total 
annual energy requirements. 

The utilization of hydroelectric power had to wait for science to 
develop methods of sending electric energy over long distances, and 
also for the building of immense dams and hydroelectric plants. 
Recently, the capacity of existing hydroelectric power plants in the 
United States totaled 25 million horsepower, while the undeveloped 
water power was estimated at 117 million horsepower. Engineers be- 
lieve that, long before we can no longer depend upon coal, petro- 
leum, natural gas, and the harnessed power of running water, nuclear 
energy will do much of the world's work. 

Motor fuels from coal. Chemists, faced with an imminent world- 
shortage of petroleum, turned to the development of methods of 
synthesizing essential liquid fuels. One of the most widely used de- 
velopments has been the changing of low-grade coal into gasoline 
and other petroleum products. The most successful results were ob- 
tained by the German chemists, Friedrich Bergius (ber'gi-obs) , Franz 



Fischer, and Hans Tropsch. In 1931, Bergius was awarded the Nobel 
prize in chemistry. 

Bergius found that coal contained half as much hydrogen as petro- 
leum. By doubling the amount of hydrogen chemically combined in 
coal, he hoped that lie could obtain a product that would resemble 
gasoline. He developed a process now called the Bergius process, in 
which a thick paste of powdered coal mixed with heavy oil is intro- 
duced into steel drums and heated to a temperature of about 400C. 
under about 200 atmospheres pressure. Hydrogen is forced into the 
mixture, which also contains a catalyst. When the hydrogenation is 
complete, a mixture of gasoline and heavier fuel oils is obtained. 
The carbon of the coal combines chemically with hydrogen, form- 
ing certain kinds of hydrocarbons, and the successful liquefaction 
of coal is achieved. 

Assuming the composition of average gasoline to be represented 
by the formula C 7 H 1(J , we may represent the hydrogenation of coal 
by the equation: 

7C+8H 2 -C 7 H 16 .-;"::;:; ' 

During World War II, the Germans demonstrated that synthetic 
gasoline from coal can be used for the same purposes as. petroleum- 
derived gasoline. In the United States, following the war> both the 
federal government and private industry conducted research on the 
production of gasoline from coal and from natural gas. It was found 
that synthetic gasoline is more expensive to produce than gasoline 
from petroleum, but would be commercially practical in the event 
of shortages of the latter. 

Apparatus in a coal-industry 
research laboratory for testing 
stoker coal and thus establish- 
ing standards for determining 
the best use for the various 

The testing laboratory of a 
coal preparation plant. The 
crucibles contain powdered 
coal which is burned in the 
electric furnace. By analysis 
of the ash, the composition 
of the coal is determined. 

National Coal Association 

Both the Bergius and Fischer-Tropsch processes were tested. The 
Fischer-Tropsch process depends either upon the reaction of steam 
and coal, or oxygen and natural gas, to form synthesis gas. 

CH-H 2 0->CO + H2 or 2CH 4 + O 2 - 2CO + 4H 2 

synthesis gas synthesis gas 

This mixture of gases is then converted with the aid of a catalyst 
into gasoline. 

7 CO + 15H 2 > C 7 H 16 4- 7H 2 O 

. catalyst , 

synthesis gas synthetic 



Bragg, William. Concerning the Nature of Things. Dover 
Publications, New York, 1954. Lecture IV on crystals covers 
the various allotropic forms of carbon. A classic. 

Storch, H. H.; Lowry, H. H.; Kiebler, M. W., Jr.; How- 
ard, H. C.; Thiessen, Gilbert; and Charmbury, H. B. "Hy- 
drogenation of Coal." Industrial and Engineering Chemistry, 
April, 1944, pp. 291-298. A round-table discussion of problems 
in this process. 


1. Destructive distillation is the heating of a complex or- 
ganic compound in a chamber from which air is excluded 
and condensing the vapors formed. 



2. Fuel value of coal is measured in calories or British 
thermal units, Btu. One Btu is the amount of heat required 
to raise 1 Ib. of water 1F. A calory is the amount of heat re- 
quired to raise 1 g. of water 1C. One Btu is equivalent to 
252 calories. 

3. Organic chemistry is the chemistry of carbon compounds. 
At present more than 650,000 compounds of carbon are known. 
Synthetic chemistry reaches its greatest development in this 
branch of chemistry and has already produced dyes, drugs, and 
solvents many of which are better than natural substances. 

4. Today, man's chief sources of energy are coal, petroleum 
and natural gas, and water power. In the years just ahead, 
nuclear energy is likely to assume greater importance. 


Group A 

1. C may act either as a metal or as a nonmetal. (a) Why? 
(b) Why is C not active at ordinary temperatures? 

2. (a) What are the three allotropic forms of C? (b) Name 
the allotropic forms of two other elements. 

3. (a) What chemical property is common to all allotropic 
forms of C? (b) What product depends on this property? 

4. What accounts for the extreme hardness of diamond? 

5. How did the action of extreme heat on diamond aid 
Lavoisier in securing acceptance of his theory of burning? 

6. (a) What are two possible products of the oxidation 
of C? (b) Write equations to illustrate. 

7. (a) What are the three chief uses of diamond? (b) Upon 
what properties do these depend? 

8. What evidence supports the accepted theory of diamond 
formation? ft 

9. (a) What are the chief uses of graphite? (b) Upon 
what properties do these depend? 

10. Describe briefly the commercial preparation of graphite. 

11. (a) What is an electric furnace? (b) Name three kinds. 

12. Make a diagram of one kind of electric furnace and ex- 
plain how it produces high temperatures. 

t . .- 

13. Make a table of four kinds of amorphous C, giving the 


method of manufacture, composition, and chief uses of each. 

14. (a) What is destructive distillation? (b) Describe the 
making of coke, and (c) name six byproducts of the process. 

15. Activated charcoal is used in gas masks. Why? 

16. Below are two lists of items. Match each item in the list 
on the left with the correct item from the list on the right. 

1) watch bearings a) activated charcoal 

2) Moissan b) graphite 

3) peat c) printer's ink 

4) Bergius d) diamond crystal 

5) lampblack e) coke 

6) manufacture of iron f) artificial diamonds 

7) lead pencils g) bacteria 

8) adsorption h) hydrogenation 

17. What weight of coke containing 80 percent C is needed 
to reduce 5 Ib. of CuO? 

T f , T , 


18. Assume the formula of gasoline to be C 7 H 16 . If 500 ml. 
of water vapor are formed during the combustion of a certain 
quantity of gasoline, what volume of CO 2 will be formed? 

19. What volume of CO 2 will be formed during the com- 
bustion oi; 2000 Ib. of coal containing 80 percent C? (Note: 
the ounce-molecular volume of any gas is 22.4 cu. ft.) 

20. According to available evidence, how were coal de- 
posits formed? 

21. Compare anthracite and bituminous coal for use in 
household heating. 

22. What factors should be considered in buying coal? 

23. How may the heat value of one coal be compared with 
the heat value of another? 

24. Complete and balance the following equations: 

a) C + H 2 -> c) 

b) Fe 2 O 3 + C -> d) SiO 2 + C -> 

25. What is meant by saying that a certain sample of coal 
liberates 14,000 Btu? 

26. What is the major use of anthracite coal? What two in- 
dustries are the major consumers of bituminous coal? 

. I . 

27. Explain the Bergius process for the hydrogenation of 


coal. Why did he experiment to achieve this result? 

28. With the aid of equations explain the synthesis of 
gasoline by the Fischer-Tropsch process. 

29. (a) What are three physical properties possessed by all 
allotropic forms of C? (b) In what physical property do they 
differ radically? 

30. Explain why the term lead pencil is a misnomer. 

31. Give two reasons why graphite is used in making stove 

Group B 

32. 2 g. of coal raised the temperature of 2000 ml. of water 
6C. Determine the fuel value of this coal in calories. 

33. Name four reducing agents and give an example of the 
reducing power of each. 

34. How is an imitation diamond distinguished from the 

35. Wooden posts are sometimes charred before being placed 
in the ground. Explain. 

36. Charcoal tablets are sometimes used in the treatment 
of certain stomach disorders. Explain. 

37. What produces chemical changes in an electric furnace: 
the heat or the electric current? Explain. 

38. How does a large percentage of ash or S in coal affect 
its quality as a fuel? 

39. (a) Why is Woehler sometimes called the father of 
organic chemistry? (b) What is the difference between organic 
and inorganic compounds? 

40. How would you identify three black powders in a mix- 
ture of graphite, manganese dioxide, and copper oxide? 


1. Place a small amount ot sugar on a hot stove or hot plate 
and report what happens, indicating as many changes as you 

2. Make a small working model of an electric lurnace used 
in the production of graphite. 

3. Large users of coal purchase coal according to definite 
specifications as to the percentage of C, volatile material, ash 
and sulfur content, and fuel value. Does your father buy coal 
according to definite specifications? Why do most domestic 
consumers not bother about specifications in buying coal? 
Is this a sensible state of affairs? Explain. 




. . . In /757 Joseph Black discov- 
ered that carbonic acid gas could not 
be breathed by animals, and had a 
poisonous effect on them. Sparrows 
introduced into an atmosphere of 
the gas died in 10 or n seconds, but 
if their nostrils were stopped with 
lard, their death took place only at 
the end of 3 or 4 minutes. Robert 
Routledge in A Popular History of 
Science, 1881 

A Scottish physician discovers carbon dioxide. Joseph Black, a 
physician, was very much interested in a lively discussion between 
two professors of medicine. One professor believed that limewater 
made from limestone was a more effective medicine than limewater 
made from oyster shells. Black, to settle this controversy in true sci- 
entific manner, undertook a thorough study of several carbonates. 
In 1754, while heating one of these carbonates (magnesium carbo- 
nate) , he obtained pure carbon dioxide gas for the first time. This 
was 20 years before the isolation of pure oxygen. Black called car- 
bon dioxide "fixed air." 

MgCO 3 - MgO + CO 2 T 

The atmosphere is a vast storehouse of this gas. Immense quanti- 
ties also occur dissolved in surface and ground water. Large quan- 
tities escape from volcanoes and crevices in the earth; it is also 
ejected in tremendous volumes from wells that produce petroleum 
or natural gas. In fact, enough carbon dioxide could be obtained 
from oil wells in the United States alone to supply the world. Carbon 
dioxide is formed when carbon or its com pounds are burned. Be- 
cause all common fuels contain carbon either free or in compounds, 



.. ~ TH. c 
oxygen cycle. 

green plants in sunlight 

carbon dioxide is produced whenever such fuels are burned. Finally, 
tremendous volumes of carbon dioxide are locked up chemically in 
the great masses of metallic carbonates and bicarbonates found in 
the earth's crust. 

Why carbon dioxide is necessary for life. Life, as we know it today, 
would be impossible without this percentage of carbon dioxide gas 
in the air, small though it is (0.04 percent) . For without it, plants 
even with the aid of sunlight could not manufacture starch. As a 
result, our food supply would diminish and finally vanish altogether. 
Since coal was formed in prehistoric days by the destructive distilla- 
tion of plants, even this great source of power would never have been 
formed if carbon dioxide had not been present in air. 

On the other hand, a much larger concentration of carbon dioxide 
in the atmosphere would be fatal, for it would dilute the oxygen and 
choke the breath of life out of living things. The delicately balanced 
ratio of oxygen to carbon dioxide in the air may be better appreciated 
after a study of the carbon dioxide-oxygen cycle. 

The carbon dioxide-oxygen cycle. Carbon dioxide is constantly 
being added to the air by the breathing of animals. It is made in the 
body tissues by the oxidation of carbon compounds in these tissues. 
Some of the oxygen inhaled in air is changed to carbon dioxide, 
which passes out of the body through the lungs. Every breath of air 
exhaled contains 100 times as much carbon dioxide as was inhaled. 
At the same time, the amount of oxygen decreases from about 2 1 per- 
cent in the air inhaled to about 16 percent in the air exhaled. 

Every ton of coal and liquid fuel burned sends about four tons of 
carbon dioxide into the air, diminishing, at the same time, the oxy- 
gen content of the air. During the decay and fermentation of or- 
ganic material, immense volumes of carbon dioxide are given off 
into the air. 

You might suspect that, after hundreds of thousands of years, 
these three processes - breathing, burning, and decay would have 
filled the air with enough carbon dioxide to destroy life. But com- 
pensating devices offset this increase in carbon dioxide. 




All the waters of the earth are continually dissolving carbon di- 
oxide, thus removing it i'rom air. In the formation of carbonates and 
bicarbonates in nature, great quantities of free carbon dioxide of the 
air are used up in the chemical weathering of rocks. Most important 
of all, during daylight hours all green plants are absorbing carbon 
dioxide, converting it into a form of sugar (fruit sugar) and into 
starch, and returning free oxygen to the air. The process that is 
responsible for this change, photosynthesis, has been described by 
some writers as the most important chemical process in the world, 
for if it should cease man and other animals would have no food. 

In higher plants, photosynthesis takes place only in parts that con- 
tain chlorophyll, a green coloring matter that acts as a catalyst. The 
final equation for the formation of starch is: 

6CO 2 + 5H 2 O 6O 2 -f C 6 H,oO 5 (starch) 

In Chapter 36 photosynthesis in relation to food-building is dis- 
cussed, but at present we are concerned particularly with its effect 
upon the composition of air and the way in which it helps to main- 
tain the balanced ratio of carbon dioxide to oxygen. This balance, 
as it. relates to plant and animal life, is indicated by Fig. 74. 

How carbon dioxide is prepared. When an acid is added to a 
carbonate or bicarbonate, carbon dioxide is liberated. In the labora- 
tory, calcium carbonate, in the form of chips of marble or limestone, 
is treated with hydrochloric acid, and the CCX is collected by water 

In these tanks, chlorophyll, the catalytic agent in photosynthesis, 
is extracted from alfalfa leaves. 

American Chlorophyll Division Strong, Cobb and Company, Inc. 


displacement, as shown in Fig. 75 below. This gas may also be col- 
lected by air displacement, for CO 2 is heavier than air. 

CaCO 3 + 2HC1 - CaCl 2 + H 2 O + CO 2 1 

Most of the carbon dioxide used commercially is obtained as a 
byproduct from coke furnaces, gas wells, and fermentation processes. 
The heating of limestone to make quicklime 

CaCO 3 - CaO + CO 2 

also furnishes carbon dioxide for industrial use. Carbon dioxide 
for industrial use is separated thoroughly from all impurities before 
being liquefied and stored in steel cylinders under pressure. It is 
known commercially as carbonic acid gas. 

Physical properties of carbon dioxide. Carbon dioxide is a color- 
less, odorless gas; it is 1| times as heavy as air, and is soluble in wa- 
ter volume for volume (at room temperature and pressure) . At 
higher pressures and lower temperatures, it dissolves in water in 
much greater volumes. 

At ordinary temperatures, a pressure of 52 atmospheres is suffi- 
cient to condense carbon dioxide to a colorless liquid. When the 
pressure over this colorless liquid is decreased, it evaporates rapidly. 
During this process, evaporation of part of the liquid withdraws heat 
so rapidly from the remaining liquid that a great amount of the 
liquid carbon dioxide is changed to solid carbon dioxide, called 
carbon dioxide snow or dry ice. 

Chemical properties of carbon dioxide. Carbon dioxide does not 
burn. It will support the combustion of only very active substances, 
such as sodium. Because it is heavier than air, it can be poured like 
a liquid. When a bottle of the gas is poured over a lighted candle, 
the flame is extinguished. Why? 

The chief chemical property of carbon dioxide is its ability to 
combine with water, forming carbonic acid. When carbon dioxide 
gas is bubbled through water, part of it combines chemically and 

CO 2 

Fig. 75. Laboratory prepara- 
tion of carbon dioxide. What 
other gat may be collected 


Fig. 76. A cubical diagram 
of the carbon dioxide mole- 




forms carbonic acid; the rest merely dissolves in the water. Carbon 
dioxide is, therefore, the anhydride of carbonic acid. 

H 2 O + CO 2 -> H 2 CO 3 (carbonic acid) 

Electron structure of carbon dioxide. Carbon dioxide is a non- 
polar, covalent compound whose structure may be represented as 
shown in Fig. 76. Notice that each oxygen atom shares two pairs of 
electrons with the carbon atom. The carbon atom shares four pairs of 
electrons with the two oxygen atoms. Since the valence of an ele- 
ment in a nonpolar compound is equal to the number of shared 
pairs of electrons, the valence of oxygen is two and the valence of 
carbon is four. 

Carbonic acid. Carbonic acid closely resembles sulfurous acid. It 
is unstable, decomposing, when warmed, into water and its anhy- 
dride, CO 2 . Thus the reaction for the formation of carbonic acid is 

H 2 C0 3 * H 2 + C0 2 

Because of the highly unstable nature of carbonic acid, it has never 
been isolated in the pure state. 

Like sulfurous acid, carbonic acid is a weak acid, dissociating only 
slightly, forming few hydrogen ions. Carbonated waters containing 
dissolved carbon dioxide are so weakly acid that they may be con- 
sumed in large quantities without ill effect. Although carbonic acid 
turns blue litmus paper pink, its sour taste is scarcely noticeable. 

Some of the many uses of carbon dioxide. The chief uses of carbon 
dioxide are: (1) in the manufacture of effervescent beverages, (2) as 
dry ice for refrigeration, (3) in the leavening of bread, (4) in extin- 
guishing fires, (5) in the manufacture of washing and baking soda, 
and (6) in the modern synthesis of urea, (NH 2 ) 2 CO, a white crystal- 
line substance used in the impregnation of wood, in the treatment 
of wounds (under the name carbamide) , and in making a valuable 
fertilizer, "Uramon": 

2NH 3 + C0 2 -> H 2 + (NH 2 ) 2 CO 

The normal human adult eliminates about 30 grams of urea as a 
waste product daily in liquid excretions. 



Rubber life rafts and life belts are inflated by means of carbon 
dioxide. The carbon dioxide so used is stored under pressure in small 
steel bulbs not unlike the carbon dioxide cartridges used on siphons 
in making charged water. 

Another use of carbon dioxide is in making car bo gen, a mixture 
of 95 percent oxygen and five percent carbon dioxide, administered to 
victims of gas poisoning and pneumonia to induce more rapid res- 
piration. Carbon dioxide in the blood appears to stimulate the re- 
spiratory nerve centers that control breathing. The heavy breathing 
of a runner, for example, is caused by the large quantity of carbon 
dioxide produced in his body during violent exercise. Adding more 
carbon dioxide to the air inhaled, therefore, causes the more rapid 
breathing needed by a patient under treatment. 

How carbon dioxide is used in effervescent beverages. Priestley 
discovered the pleasant taste of water containing dissolved carbon 
dioxide. Before the Royal Society of England he prepared "a glass 
of exceedingly pleasant sparkling water which could hardly be dis- 
tinguished from Seltzer water," and received the Society's gold medal 
for his discovery. This was the first great and unforgettable triumph 
of this amateur chemist. 

Years later, when he was forced to come to America, Priestley in- 
terested Dr. Physick, of Philadelphia, in this beverage. In 1807, 
Dr. Physick had a chemist prepare carbonated water with a little 
fruit juice for his patients. This was the beginning of the soft drink 
industry in America, which uses annually more than 200 million 
pounds of carbon dioxide. 

Soda water, so named because sodium carbonate was then used 
in preparing the carbon dioxide, is now prepared by forcing carbon 
dioxide gas into cold water at high pressures. When the pressure is 
released, the excess carbon dioxide gas is liberated, causing the bub- 
bling, or effervescing. Carbon dioxide is supplied to soda fountains 
in liquid form in steel cylinders. 

Dry ice. White, solid carbon dioxide has come into wide use for 
the refrigeration of foods especially perishable products in transit. 
It has three advantages over ice: (1) the temperature of dry ice, 
78C., is much lower than that of ice; (2) dry ic$ does not melt 
into a liquid, but changes directly into a gas; and (3) in changing 
from a solid to a gas, dry ice absorbs three times as much heat ^s ice 
does when it melts. 

Some dry ice is used in the cooling, and hardening of rivets made 
of aluminum alloys, and in shrink-fitting. This is |a process by which 
a metal fitting of correct size is expanded by hel$i|kg until it can 

Blocks of carbon dioxide dry ice, 
direct from the hydraulic presses, 
are cut into cubes before being 
wrapped for shipment. 

ill Corpnr 

be placed over or around the base to which it is to be attached. The 
fitting is then shrunk by cooling until it adheres to the base with 
great pressure. Dry ice is also used in the low-temperature drying of 
many different kinds of biological materials, in the preservation and 
shipment of blood plasma, and in the quick-free/ ing- of many dif- 
ferent kinds of foods. 

Baking powders liberate carbon dioxide lor leavening. Bread 
that is not porous is hard, unpalatable, and somewhat indigestible. 
Leavening bread makes it porous, light, and more easily digested. 
Although bread may be leavened by beating air into the dough, the 
method more generally used is to liberate large volumes of carbon 
dioxide gas in the dough by chemical action. Carbon dioxide is an 
ideal gas for leavening because it is colorless, odorless, nonpoisonous, 
and easily and inexpensively prepared. 

It is generally produced by baking powders, which are mixtures 
of two white powders, one of which is sodium bicarbonate, or baking 
soda, and the other a substance such as monocalcium phosphate, 
Ca (H,,PO 4 ) . cream of tartar, KHC 4 H 4 O,., or sodium aluminum sul- 
fate, NaAl (SO 4 ) 2 . While dry, the two powders do not react, since 
most inorganic reactions take place between ions. However, when 
water is added, the powders dissolve, dissociate, and an ionic reaction 
takes place with the liberation of carbon dioxide gas, causing the 
reaction to go to completion. One of these reactions may be repre- 
sented as follows: 

Baking soda + cream of tartar > Rochelle salt 

NaHCO 3 + KHC 4 O 6 -> KNaC,H 4 O 6 + H 2 O + CO 2 \ 

Sodium bicarbonate + potassium acid > potassium sodium 
tartrate tartrate 

Baking powders differ mainly in their speed of action. They also 
contain about 15 percent starch or flour as a filler to keep the salts 
dry and thus prevent them from reacting before being added to the 



concentrated H 2 SO 4 

Fig. 77. Cross section of a portable 

carbon dioxide fire extinguisher. concentrated 

solution of NaHCO 3 

copper tank - 

Some housewives prefer to use their own ingredients in leavening. 
For example, in making sour-milk biscuits, which many people pre- 
fer to baking-powder biscuits, the housewife uses baking soda and 
sour milk, which contains lactic acid. 

Carbon dioxide for leavening is produced also by yeast, one of 
the oldest of leavening agents. Yeast consists of living plant cells, 
which produce zymase. Zymase, which is a mixture of several enzymes, 
acts catalytically on the sugar present in the dough and breaks it 
down into alcohol and carbon dioxide (see page 548) . What advan- 
tage has yeast over commercial baking powders? 

How carbon dioxide is used in firefighting. Carbon dioxide gas is 
an ideal firefighter. It does not support combustion, is heavier than 
air, and can be quickly and cheaply liberated in large volumes. 

One kind of common portable fire extinguisher consists of a copper 
tank partly filled with a concentrated solution of sodium bicarbonate. 
Resting on a shelf inside the top of the tank is a bottle of concen- 
trated sulfuric acid covered with a loose lead stopper. When the tank 
is inverted, the sulfuric acid pours out of the bottle, and reacts with 
the sodium bicarbonate, liberating carbon dioxide gas, which carries 
out of the nozzle with it a fine spray of water and some sodium sul- 
fate. The equation is: 

2NaHCO 3 + H 2 SO 4 -> Na 2 SO 4 + 2H 2 O + 2CO 2 1 

To put out stubborn fires, such as oil conflagrations which water 
cannot extinguish, "Foamite-Firefoam" mixture is used. In this, an 
aluminum sulfate solution takes the place of the sulfuric acid in the 
ordinary type of fire extinguisher. This aluminum salt combines 
with water by a process called hydrolysis, forming gelatinous alumi- 
num hydroxide. The carbon dioxide liberated is held fast in large, 
tough bubbles by the aluminum hydroxide as well as by a sticky 
extract of licorice. Thus, the mixture spreads over the fire a layer 
of large bubbles of carbon dioxide. These smother the fire by keep- 
ing out the oxygen of the air. The equation is: 

A1 2 (SO 4 ) 3 + 6NaHCO 3 -> 2A1(OH) 3 + 6CO 2 + 3Na 2 SO 4 


Fighting fire with liquid carbon dioxide. Another form of fire- 
fighting apparatus used very widely depends on liquid carbon diox- 
ide under pressure in steel containers. The liquid carbon diox- 
ide may be in portable cylinders or in built-in systems (such as in 
airplane engines) . When the valve is opened, the liquid carbon di- 
oxide vapori/es and rushes very rapidly out of the tube. This sud- 
den evaporation of part of the carbon dioxide cools the remaining 
liquid to a white solid, and this carbon dioxide snow, played over 
the fire, quickly puts it out. The fact that carbon dioxide is able to 
penetrate obstructions without damaging equipment makes it one 
of the most rapid and efficient of fircfighting substances. 

Many ships, including the Queen Mary, are equipped to use car- 
bon dioxide for fighting fires in the various storage compartments 
throughout the ship. In a recent year, more than 500 million pounds 
of liquid carbon dioxide and carbon dioxide gas were produced in 
this country. 

Salts of carbonic acid. Because carbonic acid contains two replace- 
able hydrogen atoms, it is a dibasic acid. As you know, one or both 
of these atoms may be replaced by a metal. Hence carbonic acid 
forms two series of salts, carbonates and bicarbonates. Thus carbonic 
acid reacts with sodium hydroxide, forming sodium carbonate or 
sodium bicarbonate, depending upon the conditions. 

H 2 CO 3 + 2NaOH -> Na 2 CO 3 (sodium carbonate) + 2H 2 O 
H 2 CO 3 + NaOH -> NaHCO 3 (sodium bicarbonate) + H 2 O 

With the exception of sodium carbonate, potassium carbonate, and 
ammonium carbonate, all carbonates are insoluble in water. When a 

Demonstrating the effectiveness 
of a carbon dioxide extinguisher 
against a fire in electrical equip- 



carbonate is heated, carbon dioxide is liberated and the oxide of 
the metal remains. 

CaCO 3 - CaO + CO 2 f 

The carbonates of sodium and calcium are discussed in Chapters 30 
and 31. 

How to test for a carbonate or bicarbonate. An acid liberates 
carbon dioxide gas from either a carbonate or a bicarbonate. If this 
gas is then passed into limewater, Ca (OH) 2 , a white precipitate, cal- 
cium carbonate, CaCO 3 , forms. 

Ca(OH) 2 + C0 2 -> CaC0 3 1 + H 2 O 

On bubbling more carbon dioxide gas through the mixture, the 
white precipitate disappears as a result of the solubility of calcium 
bicarbonate which is formed. 

CaCOg + H 2 O + CO 2 -> Ca(HCO 3 ) 2 (calcium bicarbonate) 


Armstrong, George B. "Dry Ice." Chemistry, Feb. 1951, 
pp. 24-31. An excellent, illustrated article on the history, 
properties, manufacture of and uses of solid carbon dioxide. 

U.S. Department of Agriculture, 1945. Safe Use and Storage 
of Gasoline and Kerosene. Farmers' Bulletin No. 1678. Supt. of 
Documents, Washington, D.C. 10^. Some excellent advice on 
putting out gasoline and kerosene fires. 




and . 



Group A 

What are the chief sources of CO 2 ? 

How and by whom was CO 2 discovered? 

In what way does life depend upon CO 2 ? 

Copy and complete the following statements. Do not 

in this book. CO 2 is being added to the air constantly 

. , . . . , and .... CO 2 is removed from the air by . . . , . . . , 

Make a diagram representing the carbon dioxide-oxygen 
By what three commercial methods is CO 2 produced? 


7. Write the equation for one commercial method of pro- 
ducing CO 2 . 

8. Make a labeled diagram for the laboratory preparation 
of CO 2 . 

9. (a) Write the chemical equation for the laboratory 
preparation of CO 2 . (b) Why does this reaction go to com- 

10. (a) Can carbon dioxide be collected by air displace- 
ment? (b) Explain. 

11. What are the physical properties of CO 2 ? 

12. What are three chemical properties of CO 2 ? 

13. (a) Write the reversible reaction of carbon dioxide and 
water, (b) How could you make large quantities of CO 2 com- 
bine with H 2 O? (c) The behavior of CO 2 with water places it 
in what class of compounds? 

14. What other acid closely resembles H 2 CO 3 ? 

15. (a) What are six uses of CO 2 ? (b) Opposite each use 
state the property or properties on which that use depends. 

16. Describe briefly the beginning of the soft-drink industry 
in America. 

17. (a) How is dry ice made? (b) Dry ice has what advan- 
tages over ice? t 


18. What is the general composition of baking powder? 

19. What is the function of each ingredient in one type of 
baking powder? 

20. Write a balanced equation representing the liberation 
of carbon dioxide from one type of baking powder. 

21. (a) Describe the chemical action of some other type of 
leavening agent, (b) What advantage has it over commercial 
baking powders? 

22. What are four reasons that CO 2 is an ideal leavening 

23. A manufacturer makes enough baking powder to pro- 
duce 15 tons of CO 2 . What weight of NaHCO 3 does he use? 

24. A sample of baking powder liberated 200 ml. of CO 2 . 
Considering that the reaction went to completion, what weight 
of NaHCO 3 was used if the baking powder was of the cream- 
of- tartar type? 



25. What volume of air would be needed to burn completely 
enough coke to make 10,000 cu. ft. of CO 2 ? 

26. Why do baking powder reactions go to completion? 

27. Describe with diagram and equation the principle of 
one type of portable fire extinguisher. 

28. (a) H 2 O is not a good agent for putting out a fire of 
burning oil. Why? (b) Why is CO 2 from the ordinary type of 
fire extinguisher ineffective? (c) What principle is used in the 
"Foamite-Firefoam" method? 

29. (a) H 2 CO 3 has two series of salts. What are they? 
(b) Write two equations to illustrate. 

30. What is the action of heat on carbonates? 

31. (a) Describe a complete test for a carbonate, (b) Write 
the equation for this chemical test. 

32. When sour milk is used in cooking, NaHCO 3 is used 
instead of, or in addition to, baking powder. Explain. 

33. How can CO 2 be removed from a sample of air? 

34. Can pure GO 2 alone be used in fighting fire? Explain. 

Group B 

35. (a) Why does CO 2 collect in wells and caves? (b) What 
is its source? 

36. Does the effervescence of a solution when an acid is 
added prove the liberation of CO 2 ? Explain. 

37. 250 ml. of a gas weigh 0.49 g. This gas contains 27.3 
percent carbon and 72.7 percent oxygen. Find the formula of 
this compound. 

38. Most chemical reactions take place between ions. How- 
ever, CaCO 3 is insoluble in water. Explain the liberation of 
CO 2 when HC1 is added to solid CaCO 3 . 

39. Explain why a knowledge of chemical arithmetic is es- 
sential in preparing baking powders. 

40. SO 2 is a reducing agent but CO 2 is not. Explain. 


1. Make a study of the cost per ounce of a half-dozen widely 
advertised baking powders. Nearly all baking powders liberate 


the same amount of CO 2 for equal weights of baking powder. 
Is the wide range of prices for such baking powders warranted? 
Give reasons for your answers. Why do such wide ranges in 
prices exist? 

2. Question your mother, or other members of your family, 
and friends about their preferences in leavening agents. List 
the data you obtain and give your own reactions to such pref- 

3. Visit a soft-drink bottling plant in your neighborhood, 
and report on the method used in charging the soft drinks 
with CO 2 . 

4. After consulting your teacher of biology on sources of 
information, prepare a two- or three-page report on micro- 
organisms, other than yeast, that are used by man in the prep- 
aration of his food supply. Arrange a class discussion on this 

5. Construct a glass model of an acid-sodium bicarbonate 
type of portable fire extinguisher. Demonstrate its action be- 
fore the class. 



. . . The introduction of gas light- 
ing (1810) met with much opposi- 
tion, as one can understand, on the 
part of the tallow chandlers and sell- 
ers of whale oil at that time used 
as an illuminant; but after a time 
the new method of lighting was wel- 
comed. Alexander Findlay, 1930 

Priestley discovers an oxide of carbon that burns. Almost half a 
century after the discovery of carbon dioxide, another quite differ- 
ent oxide of carbon was found, carbon monoxide, CO. Not until the 
last year of the eighteenth century was carbon monoxide known to 
be a new compound of carbon. 

Priestley again was the first to recognize clearly a new gaseous 
compound in the flickering blue flame which played over his fur- 
nace fire. This valuable discovery was made in America, where this 
dissenting English minister had fled to escape political and religious 

How carbon monoxide is formed. Priestley's observation indicated 
that when coal or other carbonaceous fuel is burned in a limited sup- 
ply of air, carbon monoxide, rather than carbon dioxide, is formed. 

2C + O 2 -> 2CO 

Carbon monoxide is formed also in a house furnace by the reduc- 
tion of carbon dioxide gas as it passes over glowing coal, which, as 
you know, is an excellent reducing agent. 

CO 2 + C - SCO 


Normally, carbon monoxide burns at the surface of a coal fire, 
giving off heat and changing to carbon dioxide, but if the supply of 
air is insufficient, or if the flame is chilled, some of the carbon mon- 
oxide may escape without burning. 

How carbon monoxide is prepared in the laboratory. Formic acid 
is a colorless liquid with a characteristic odor, that is, a distinct odor 
that can be used to identify it. Its formula is HCOOH (only the 
last H is a replaceable hydrogen atom) . When warm concentrated 
sulfuric acid, a powerful dehydrating agent, is added to formic acid, 
it extracts the HOH of the acid, leaving pure CO. 

HCOOH -> HOH + CO t 

Physical properties of carbon monoxide. Carbon monoxide is a 
colorless, tasteless, almost odorless gas, slightly lighter than air, and 
slightly soluble in water. Its odor is very difficult to detect. It can be 
liquefied, but only under high pressures and low temperatures. 

Chemical properties of carbon monoxide. Carbon monoxide burns 
in air with a pale blue flame, forming carbon dioxide, which, because 
it is unable to unite with any more oxygen, cannot burn. 

2CO + 2 - 2C0 2 

Because of carbon monoxide's great attraction for oxygen, it is 
an excellent reducing agent, and therefore it is used in the extraction 
of iron from its oxide. The carbon monoxide is formed by the reac- 
tion of coke (carbon) and CO 2 in the upper part of a blast furnace. 

Fe 2 O 3 + SCO -> 2Fe + 3CO 2 T 

Carbon monoxide combines with certain metals, forming a series 
of compounds called carbonyls. For example, when nickel is heated 
in the presence of carbon monoxide to a temperature of about 
40C., nickel carbonyl, a poisonous liquid, is formed. 

Ni + 4CO -> Ni(CO) 4 (nickel carbonyl) 

Upon further heating, carbonyls liberate CO and give up their met- 
als. Nickel carbonyl is used in the extraction of nickel from its ores. 
Carbon monoxide in the presence of suitable catalysts also com- 
bines directly with hydrogen, forming wood alcohol (see page 548) , 
and with chlorine, forming phosgene. 

CO + 2H 2 - CH 3 OH (wood alcohol) 
CO + C1 2 > COC1 2 (phosgene) 


Some uses of carbon monoxide. Carbon monoxide is an excellent 
fuel, burning with intense heat and leaving no solid residue. It is 
most often used in conjunction with other gases as a mixed fuel, as 
in water gas, for example, which is a mixture of carbon monoxide 
and hydrogen. 

Carbon monoxide is used widely in the metallurgy (extracting a 
substance from its ores) of iron and nickel. 

Deadly carbon monoxide, a cumulative poison. Carbon monoxide 
is the most widespread poison connected with human life and activ- 
ity. It acts painlessly. Furthermore, it gives no warning of danger, 
because it is almost odorless, and its victim slowly passes into a state 
of unconsciousness. Its physiological action is caused by the strong 
affinity of carbon monoxide for hemoglobin, an attraction that is 300 
times greater than the attraction of oxygen for hemoglobin. 

Hemoglobin is present in the red blood corpuscles and its function 
is to unite chemically with oxygen and carry it to all parts of the body 
by the circulation of the blood. By forming a stable compound with 
the hemoglobin in blood, carbon monoxide prevents oxygen from 
reaching millions of cells in the body. The victim is thus killed by 
suffocation a lack of oxygen. 

The formation of carbon-monoxide-hemoglobin takes place gradu- 
ally, and this substance remains in the blood. When the percentage 
of this stable compound has reached 40 percent, the victim collapses; 
when it reaches 80 percent, death occurs. Such a poison, which col- 
lects slowly in the body and over a long period of time until it reaches 
a dangerous concentration, is said to be a cumulative poison. 

The treatment for carbon monoxide poisoning consists of induc- 
ing respiration by the Schafer prone-pressure or the more recently 
adopted Holgar-Nielsen "back-pressure, arm-lift" method in the open 
air, while carbogen is administered. No alcohol should be given. As 
you know, the carbon dioxide makes the patient breathe more deeply. 
The injection into the blood stream of methylene blue, a dye, as an 
antidote for carbon monoxide poisoning has been reported to be 
effective in certain cases. 

Where the danger of carbon monoxide poisoning lurks. The ex- 
haust gas of automobiles, in which carbon monoxide exists in vary- 
ing percentages up to about 12 percent, is one of the most common 
sources of carbon monoxide. Hundreds of persons are killed by this 
gas every year, usually in some such ways as these. 

1) A motorist drives his car into his garage. It is cold, and he shuts 
the doors of his garage to keep warm. He keeps the motor running 
while he gets under the machine to make some repairs. Before he is 



aware of it, carbon monoxide from the exhaust pipe of the automo- 
bile causes unconsciousness. Death soon follows. 

2) On a very cold morning a driver may start the engine before 
opening the garage door, or he may sit in a closed car with the en- 
gine running while waiting for a friend. Carbon monoxide gas from 
the exhaust finds its way into the confined space and takes its toll. 

The carbon monoxide comes from the incomplete combustion of 
the gasoline. Although gasoline is a mixture of hydrocarbons, we may 
represent its composition by the formula C 7 H 16 . During incomplete 
combustion, the following reaction is one that may take place: 

C 7 H ]6 + 8O 2 - 6CO + 8H 2 O + CO 2 

The most serious industrial cases of carbon monoxide poisoning 
occur from exposure to gases, such as producer and blast furnace 
gases used in many manufacturing plants (see pages 359 and 407) . 

Most poisonings from carbon monoxide outside of industrial and 
auto cases are caused by the escape of gas from kitchen ranges and 
stoves, gas refrigerators and other gas appliances because of leaks 
and the accidental extinguishing of the burning gases. A furnace, im- 
properly adjusted, may give off quantities of this gas and become a 
menace to life. As shown below, every coal stove and furnace should 
have an adequate circulation of air, otherwise carbon monoxide may 
be formed and finally reach dangerous concentrations. 

Firing for complete combustion. In firing a coal-burning furnace, 
the object is to obtain as much heat from the coal as possible. This 
can be done only when combustion of the coal is complete, or nearly 
so, with a minimum of intermediate products, including carbon 
monoxide and free carbon, or soot, remaining unburned. Incom- 
plete combustion is dangerous because of the dangers from the car- 
bon monoxide produced. It is also wastefully inefficient and robs 



Fig. 78. Stove in operation. 
When damper A is open, but 
dampers B and C are not, in- 
complete combustion results. 
When and C only are open, 
the fire burns slowly. 

2C0 2 ' 



burning CO 
reduction of CO 2 

of coal 



the consumer of many Btu per dollar of fuel. In addition, the great 
quantities of smoke and soot produced are a menace to health and 
greatly increase expenditures for laundry and dry cleaning. 

Complete combustion of coal depends upon adequate quantities of 
air and upon proper firing methods, that is, proper placement of coal 
in the firebox. In adding coal to a burning fire, push the glowing coals 
from the front of the firebox to the rear, and place the additional coal 
in the space from which the coals were pushed. In this way, the fire is 
not smothered and great clouds of smoke do not result. The glowing 
coals at the rear of the firebox heat the added coal, combustible va- 
pors are formed and oxidized relatively completely in the region sep- 
arating the new coal from the glowing coals. 

If the new coal is spread completely over the glowing coals the 
fire is partially smothered. Combustible vapors are formed, to be 
sure, but for a long period, the temperature above the added fuel is 
lower than the kindling temperature of the vapors, and very small 
amounts of heat are produced. 

Proper regulation of the air supply is a most important factor in 
securing complete, or nearly complete, combustion. If the supply of 
air is insufficient, incomplete combustion results, and carbon mon- 
oxide and free carbon, or soot, are produced in large quantities. If 
too much air is available, combustion is relatively complete but very 
rapid, and results in uneven heating. If far too much air is available, 
much of the heat produced is carried out the chimney without trans- 
ferring its heat to the water, steam, or air in the heating system. 

Finding the damper settings that result in relatively complete com- 
bustion is not difficult, but it does require some patience and much 
observation, for under varying outside conditions, damper settings 
vary also. If you can adjust the dampers so that the flame produced 
in the firebox has a minimum of yellow flame, the conditions neces- 
sary for nearly complete combustion probably are present. 

Coal fires Should not be poked or shaken as often as most persons 
poke and shake them. Shaking of ashes should occur only when the 
accumulation of ashes becomes so deep that the air supply may be im- 
peded. Ashes insulate the metal grates of the firebox from the in- 
tense heat of the fire, and if they are all removed, the grates may be 
damaged as a result of receiving more heat than they were designed 
to receive. Poking a fire frequently spreads the burning coals through- 
out the firebox and results in too rapid combustion^ 

The simple principles just discussed will, if followed, do much 
to increase the heat that consumers get from the c6al they buy; smoke 
nuisances will be reduced, and fuel bills will be ctil itiaterially. 



How carbon monoxide can be removed from air. For a long time 
no simple method of removing noxious carbon monoxide from air 
was known. Activated charcoal, which adsorbs various gases, such as 
ammonia, acetone, benzene, and chlorine, will not remove carbon 
monoxide from air. 

A catalyst called "Hopcalite," which consists of a mixture of me- 
tallic oxides, causes carbon monoxide to change to carbon dioxide 
at ordinary temperatures and pressures. In the presence of "Hop- 
calite," 1 molecule of carbon monoxide combines with another atom 
of oxygen from the air. A rescue-breathing apparatus containing a 
"Hopcalite" canister is now used by persons who find it necessary 
to enter regions where concentration of carbon monoxide is high. 

How can carbon monoxide be detected? Simple color-detector 
tubes have been devised for determining minute amounts of carbon 
monoxide in air. One such preparation called hoolamite contains 
specially prepared iodine pentoxide, L() r> , which oxidizes carbon 
monoxide to CO 2 . 

5CO + I 2 O& - 5CO 2 + I 2 

The iodine fumes liberated cause a change in color that is directly 
proportional to the amount of carbon monoxide present. This 
amount can be determined by comparison with a standard color 
scale. The color fades, and the tubes may be used again. Before this 
method was developed, canaries were often used as detectors because 
these birds are very sensitive to minute amounts of this poison and 
show its effects before man does. 

Official U.S. Navy photograph 

Navy firefighters ready for 
action. Their rescue-breathing 
apparatus will protect them 
from carbon monoxide and 
other dangerous fumes. 

These recording devices keep 
a constant check on the con- 
dition of the air in the Lincoln 
Tunnel. The 8000-foot vehi- 
cular tunnel connects the 
New York and New Jersey 
banks of the Hudson River. 

The Port of New York Authority 

In some tunnels, such as the Holland and Lincoln Tunnels at 
New York City, machines have been installed which record on a 
time chart the amount of carbon monoxide present in the air. A 
warning bell is caused to ring when the carbon monoxide reaches 
a concentration of four parts in 10,000. At this point, also, dampers 
and fans begin to operate automatically to change the air. 

What are the more important gaseous fuels? The fuel gases 
burned each year in this country are worth more than two billion 
dollars. The total volume of gas produced is about eight trillion 
cubic feet. Besides pure carbon monoxide, which is seldom used 
alone, the most important gaseous fuels are water (or synthesis) gas, 
producer gas, coal gas, natural gas, acetylene gas, and hydrogen. Hy- 
drogen has already been discussed; a brief discussion of the others 

A gas from steam and carbon. Water gas is manufactured by pass- 
ing water in the form of steam over glowing coke or hard coal. The 
coke is first burned in a draft of air until it is red-hot. The air is 
then shut off and the steam turned on. The temperature of the coke 
falls gradually because the reaction is endothermic, and when it 
reaches about 1000C., the steam is cut off and the air supply re- 
newed. During the process the carbon, which is an ideal reducing 
agent, combines with the oxygen of the water, leaving hydrogen. 

H 2 O + C -> CO + H 2 

water gas 

At a temperature of about 1000C., the most gas is produced. 
Above this temperature, the carbon monoxide formed may react 
with the steam and change to carbon dioxide. 

H 2 O + CO -> CO 2 + H 2 



Water gas is thus a mixture of carbon monoxide and hydrogen. 
It is used, either alone or mixed with coal gas, for domestic heating 
purposes. Since both carbon monoxide and hydrogen burn with al- 
most colorless flames, a mixture of the two cannot be used for illumi- 
nating purposes unless it is made luminous by injecting into it gas- 
eous hydrocarbons from petroleum, which, on burning, give a yellow 
flame. This process of adding hydrocarbon vapors to water gas is 
called enriching the gas. When water gas burns, it forms water va- 
por and carbon dioxide. 

CO + H 2 + O 2 - H 2 O + CO 2 T 

water gas 

When used in the synthesis of gasoline and other chemicals it is 
called synthesis gas. 

The gasification of coal. In areas where natural gas is not available 
as a fuel for factories and homes, producer gas is most generally 
used. It is produced by burning low-grade coal in a furnace with a 
limited supply of air. The chief product formed during the incom- 
plete combustion of the coal is carbon monoxide. The gas issuing 
from this furnace is mixed with large quantities of nitrogen, which 
is too inactive to unite with the coal. Producer gas, then, is chiefly 
a mixture of about 60 percent nitrogen and 30 percent carbon mon- 
oxide. It contains also about ten percent hydrogen. 

2C + air (0 2 + N 2 ) - 2CO + N 2 

Producer gas is a much poorer fuel than water gas because it con- 
tains such a large amount of nitrogen, an incombustible gas. Its 
manufacture constitutes the most efficient means of converting low- 
grade coal into power. 

A gas similar in composition to producer gas was first proposed by 
Mendeleyeff and later produced in 1933 in the Soviet Union by 
burning coal underground. Instead of mining the coal in the usual 
way, a coal seam was sealed off, set on fire, and the gases produced 
were brought to the surface through pipes. Later, pure oxygen was 
mixed with the blast of air sent into the mine, and the gas that came 
out of the mine was richer in composition than producer gas. 

This gasification of coal underground (gas mining) , held feasible 
by the U.S. Bureau of Mines after several years of experimentation, 
may turn out to be a revolutionary development. By this method we 
may be able to recover valuable coal in mines abandoned because of 
the low-grade nature of the coal or because of the thinness of the coal 


Fig. 79. Underground gasiflcc 
tion of coal. 




vein. In addition, this method makes coal working safer and health- 
ier, a consideration which is always desirable. Another advantage is 
pipeline distribution of this fuel. 

Natural gas, the Cinderella fuel. More than 80 percent of the 
gaseous fuel consumed each year in the United States is natural gas. 
In various sections of the world, especially where coal and petroleum 
deposits are found, natural gas issues from the earth when porous 
rocks saturated with it are tapped. The gas may flow out under pres- 
sure, or pumping may be required. 

Many gas wells do not yield oil, but an oil well almost always 
produces both gas and oil. The gas from wells that do not yield oil is 
usually very rich in methane, CH 4 , some wells yielding as high as 
95 percent methane. 

In the early days of the petroleum industry, little or no use was 
found for natural gas, and most of the wells were ignited and allowed 
to burn for years. When the waste of natural resources involved in 
such practices was realized, controls were established and the burn- 
ing of natural gas greatly diminished. 

Methane, also called marsh gas, is a colorless, practically odorless, 
insoluble gas which burns with an almost colorless flame, forming 
CO 2 and water vapor. 

CH 4 + 2O 2 -> CO 2 + 2H 2 O 

It has a calorific value about twice that of manufactured gas which 
it has largely replaced. It is about half as heavy as air. Natural gas is 
purified, some high quality casing-head gasoline being obtained in 



this process. H 2 S is also removed before it is sent into pipelines. It is 
also stripped of its heavier hydrocarbons such as butane and propane 
before being piped many hundreds of miles to supply factories, farms, 
and homes with light, heat, and power. For example, there is an 
1840-mile pipeline (the Big Inch) from Texas to New York. 

Probably the origin of our great supply of natural gas is the same 
as that of petroleum, which most scientists believe to be the result of 
the incomplete decomposition of vegetable or animal matter, either 
with or without bacterial action. The formation ,of petroleum does 
not require millions of years as was formerly believed. This is demon- 
strated by the formation at the present time of petroleum in off-shore 
marine sediments. Some scientists believe, however, that natural gas 
originates from the interaction of metallic carbides and water, just as 
acetylene is formed by the reaction between calcium carbide and 

CaC 2 + 2H 2 - Ca(OH) 2 + C^ t (acetylene) 

A gas from a gray solid. The gaseous fuel called acetylene, C 2 H 2 , 
is colorless and odorless when pure, very slightly soluble in water, 
and somewhat toxic. Acetylene has a tendency to explode when lique- 
fied. The gas is therefore not liquefied, but is forced at low pressure 
into a solvent called acetone, a colorless liquid obtained from the 
destructive distillation of wood, and mixed with some inert porous 
material, such as wood charcoal or asbestos fiber. It is sold in port- 
able steel cylinders and is used widely in oxyacetylene torches. When 
the valve of one of these cylinders is opened, the pressure is reduced, 
and some of the gas escapes from solution. 

For emergency use, and in lighthouses and isolated districts where 
electric lighting and illuminating gas are not available, special acety- 
lene generators have been constructed. These special generators al- 
low water to come in contact with calcium carbide at regulated rates 
of speed, so that the gas may be liberated as needed. 



CaC 2 


Fig. 80. (left) Carbide-to-water acetylene generator for large installations, (center) 
Water-to-carbide generator, (right) The acetylene burner. 

The burner used with acetylene gas must be specially constructed 
to permit the access o a large amount of air to the burning gas. 
Because acetylene is rich in carbon, it would otherwise burn incom- 
pletely and produce a smoky flame. 

2C 2 H 2 + 5O 2 > 4CO 2 + 2H 2 O (complete combustion) 
2Q>H 2 + O 2 > 4C -h 2H 2 O (incomplete combustion) 

Alcohol, acetone, vinegar, chloroform, plastics, and synthetic fibers, 
rubber, and gasoline have been built up chemically from acetylene 
gas (see Chapters 34 and 35) . 

A gas from the destructive distillation of coal. About the time 
that Priestley was studying carbon monoxide, coal gas was intro- 
duced as an illuminant. William Murdock, a Scottish workman em- 
ployed by James Watt, developer of the steam engine, carried out 
experiments that led to lighting part of the Boulton and Watt fac- 
tory in Birmingham, England, with gas in 1798. A few years later, 
gaslighting was introduced in the United States. Much opposition 
was raised against it, but the advance of science could not be stopped 
for long. Manufacturers of candles and whale oil, fought against il- 
luminating gas, even as other manufacturers have since struggled un- 
successfully against other innovations. 

Coal gas is obtained from coal by destructive distillation. Fig. 81 
shows the method of manufacture and purification in a byproduct 
coke oven. Coal is heated in a closed oven. The vapors formed are 
first passed into the hydraulic main, where some of the vapors of 

impure gas 

Fig. 81. Steps in the production of coal gas. 




coal tar and ammonia are condensed. The remaining vapors then 
enter the condensers, where the rest of the ammonia is absorbed and 
the coal tar condensed. The purifier, containing iron oxide or lime, 
removes hydrogen sulfide and other sulfur compounds. The purified 
coal gas then enters the gas holder. Coal gas contains about ten per- 
cent carbon monoxide, 40 percent hydrogen, and 40 percent methane. 
It contains also about ten percent nitrogen. 

The destructive distillation of one ton of coal yields approxi- 
mately 10,000 cubic feet of coal gas, 20 gallons of ammonia water, 
and 120 pounds of coal tar. About 1400 pounds of coke remain in 
the retort. A tarry matter called pitch is left also, which is used as a 
binder in road construction. 

What is an explosive mixture? All the gases discussed in this 
chapter burn quietly in air. However, when they are mixed with air 
in the right proportions to secure nearly complete combination and 
then ignited, very rapid oxidation takes place, suddenly producing 
extremely large volumes of gases. The high temperature of the re- 
action helps to account for the large volume of gases formed, since 
gases expand as their temperature is raised. High temperatures also 
result in high pressures. If the explosive mixture is confined, as, for 
example, in a mine or factory, great destruction takes place. 

Gases and vapors differ from one another in the range of composi- 
tion of their explosive mixtures. Thus coal gas will explode when 
anywhere from six percent to 29 percent of it is mixed with air. Air 
containing less than six percent or more than 29 percent of coal gas 
will not explode. 

The character of the explosive mixture is of prime importance in 
the working of internal-combustion engines. For example, the degree 
of smoothness and of power in the running of an automobile engine 
depends upon getting the right mixture of air and gaseous fuel 
admitted to the cylinders and ignited at the proper instant by an 
electric spark. The closer the proportions of the mixture to those 
necessary for securing complete combustion, the greater the power 
produced and the less the waste of fuel. 

Because of the importance attached to securing this proper mix- 
ture, the carburetor, in which the mixing is done, is often called the 
heart of the internal-combustion engine. This needs adjustment from 
time to time to insure highest efficiency (see Fig. 82) . 

Gasoline vapor mixed with air may form a very explosive mixture. 
Many serious accidents have resulted from the careless use of gaso- 
line in dry cleaning at home. Such cleaning operations should be 
done in the open and away from any source of ignition. In using 



mixture of 
gas and air 



spark plug 

Fig. 82. Carburetor and cyl- 
inder assembly. The gaso- 
line-air mixture is ignited in 
*he cylinder by the spark 

P |ua - The ner ay of tne ex - 

plosion is transformed into 
motion by the piston. 

gasoline for dry cleaning, vigorous or continued rubbing of the fab- 
ric should be avoided lest the friction ignite the gasoline. 

Safety measures against mine explosions. For centuries the igni- 
tion of explosive mixtures, especially in coal mines, has caused seri- 
ous loss of life. In 1556, Agricola (-grik'6-la) published a book on 
mining, one section of which dealt with "the ailments and accidents 
of miners and the methods by which they can guard against these." 
But life was cheap in those days and little was done to protect min- 
ers against explosions. At the end of the eighteenth century, there 
came an emphasis on the rights of man and with it a new humani- 
tarianism. An interest in occupational accidents and diseases was 
aroused. Gases in mines and mine ventilation were studied. 

Davy, who discovered sodium and potassium, and made other im- 
portant contributions to chemistry, was among those interested in 
the plight of the miners in England. He devised a simple safety lamp 
to prevent mine explosions. It is based upon two principles: (1) An 
explosive mixture does not undergo chemical change until its kin- 
dling temperature is reached. (2) Metal surfaces spread heat rapidly. 

In the Davy lamp, the flame is surrounded by a wire gauze, which 
distributes the heat produced by the flame over a wide area, and thus 
prevents the explosive mixture outside the lamp from reaching its 
kindling temperature. A lighted match brought over a wire gauze, 
as shown in the illustration below, will set fire to the gas above the 
gauze, but the gas below will not catch fire. Why? This is an illustra- 
tion of the principle of the Davy safety lamp. 

In commercial mining, the Davy lamp has been replaced by a 
battery-operated electric lamp. 

gas flame 

wire gauze 

gas below its 
kindling temperature 

Fig. 83. Demonstration of the 
principle of the Davy lamp. 
Why does the gas below the 
wire gauze not catch fire? 

In an underground coal 
mine, the air is tested for 
the presence of methane 
immediately after coal is 
blasted from the seam. 
The flame of the safety 
lamp turns blue when the 
gas is present. 

National Coal Association 


Faraday, Michael. The Chemical History of a Candle. 
"Kings' Treasuries of Literature" Series, E. P. Button 8c Co., 
New York, 1920. This book consists of six lectures delivered 
by Faraday before young boys and girls at the Royal Institution 
of London in 18(>()-1861. Lecture I covers (lames sources and 
structure. The book is a classic, and though about 100 years 
old may still be read with pleasure and profit. 

Manchester, Harland. New World of Machines, pp. 174- 
189. Random House, New York, 1945. "Power for Tomorrow" 
is a fine chapter in this carefully written book. 

Oettingen, W. F. von. Carbon Monoxide: Its Hazards and 
the Mechanism of Its Action. Public Health Bull. No. 290, 
1944. Supt. of Documents, Washington, D.C. 35^. 

Group A 

1. Where and by whom was pure GO first studied? 

2. (a) How is CO formed in a furnace? (b) Write an equa- 
tion to illustrate, (c) Write the equation for the oxidation of 

3. (a) How is CO usually prepared for laboratory use? 
(b) What is the function of the H.,SO 4 used? 

4. Compare the physical properties of N 2 and CO. 

5. (a) What are two chemical properties of CO in addition 
to its combustibility? (b) Write equations to illustrate each. 

6. Write balanced equations for the following. Name the 
products, (a) Action of chlorine and carbon monoxide, 
(b) action of hydrogen and carbon monoxide, (c) action ol 
carbon monoxide on heated nickel, and (d) formation of 



, , | , t T 

7. Discuss four uses of CO. 

8. (a) Explain how CO acts on the hemoglobin of the 
blood of a person breathing it. (b) What first-aid treatment 
should be given a person who has been overcome by CO? 
(c) What is the function of the small amount of CO 2 ad- 

9. Poisoning by CO often occurs in closed garages, 
(a) Why? (b) What precautions should be taken to prevent 
this danger? 

10. (a) What gases are generally adsorbed by the C used in 
gas masks? (b) What substance is used in CO detectors? 

11. (a) What are two important factors in firing with coal 
for complete combustion? (b) What is the most efficient 
method of firing? (c) Why? 

12. In firing with coal, what are the results of (a) inade- 
quate air supply? (b) too much air? (c) far too much air? 

13. How can you tell when the conditions necessary for 
complete, or nearly complete, combustion are likely to be 

14. (a) What do large quantities of black smoke issuing 
from a chimney indicate? (b) Why? 

15. (a) Is black smoke undesirable? (b) Why? 

16. (a) What are two pure gaseous fuels? (b) What are 
three that are mixtures of gases? (c) Which of (b) contain 
large quantities of CO? 

17. (a) Write an equation for the manufacture of water 
gas. (b) What is "synthesis gas"? 

18. What volume of steam is used in making 1500 cu. ft. of 
water gas? 

19. What weight of coal, containing 85 percent C, is used 
in making 2000 cu. ft. of water gas? (One , ounce-molecular 
weight of a gaseous substance occupies 22.4 cu. ft.) 

20. Water gas contains 60 percent CO and 40 percent H 2 . 
What volume of air is necessary for tlie complete combustion 
of 200 cu. ft. of this gas? 

21. (a) What is meant by enriching water gas? (b) Why 
is it done? 


22. Producer gas contains about 60 percent noncombustible 
N 2 , yet it is used in great quantities as a fuel for gas engines, 
(a) How does this N 2 affect its fuel value? (b) What process 
uses producer gas? Why? 

23. Explain the gasification of coal underground. 

24. (a) On the basis of present evidence, what do we think 
was the source of natural gas? (b) How does its present con- 
sumption rank with that of other gaseous fuels? (c) What is 
its chief constituent? 

25. (a) What are the properties of methane? (b) What 
other name has it? (c) Why? 

26. Make a table showing the source, composition, and chief 
use of four gaseous fuels. 

27. Describe briefly the stages in the manufacture of puri- 
fied coal gas. 

28. (a) What is destructive distillation? (b) Name the chief 
products of destructive distillation of bituminous coal? 

29. (a) How is C 2 H 2 made? (b) What are its principal 
properties? (c) Why must a C 2 H 2 burner be so constructed 
that it allows the access of large quantities of air? 

30. How many cubic feet of air are used in burning com- 
pletely 2500 cu. ft. of C 2 H 2 ? 

31. How would you identify CO 2 , H 2 , CO, N 2 , and NO? 

32. (a) What happens when an explosive mixture is ignited? 
(b) Why are explosive mixtures dangerous? (c) What is the 
most important factor to consider in the mixing of gases for 
explosive effect? 

33. (a) How is the explosive mixture in an automobile 
regulated? (b) Why? 

34- Describe a situation at home in which a dangerous ex- 
plosive mixture of gases might be formed. 

35. What might happen if the ventilating system in an 
underwater tunnel such as the Lincoln Tunnel in New York 
City suddenly got out of order? 

Group B 

36. Why does CO burn, whereas CO 2 does not burn? 

37. In making water gas, what would happen if the tempera- 
ture used were too high? 

38. Why has the byproduct furnace replaced the beehive 
coke oven? 


39. Why is a blue flame seen when fresh coal is added to 
burning fuel? 

40. A gas sample from a sealed fire area in a mine shows 
CO, four percent; O 2 , ten percent; CH 4 , seven percent; and N 2 , 
79 percent. Is the fire bla/ing, or is a methane explosion pos- 
sible? Explain. 

41. Suppose a sample of gasoline is half hexane, C 6 H 14 , and 
half heptane, C 7 H 16 . How many cubic feet of air are necessary 
for the complete combustion of 20 cu. ft. of this gasoline 


1. If you live near an oil or natural-gas field, visit the gas 
wells and write a report on the gas waste, if any. Do you think 
Federal control of such waste desirable? Explain. 

2. Davy refused to take out a patent on his miner's lamp, 
saying, "No, my good friend, I never thought of such a thing; 
my sole object was to serve the cause of humanity, and it I have 
succeeded I am amply rewarded in the gratifying reflection of 
having done so." Write a report expressing your views on this 
incident. Cite similar instances. 

3. Study the recently introduced Holgar-Nielsen "back-pres- 
sure, arm-lift" method and compare it with the Schater prone- 
pressure method. 

4. Using a test tube, a one-hole rubber stopper, a short de- 
livery tube and some dry sawdust or wooden splint, show that 
a combustible gas can be obtained Irom the wood similar to 
coal gas from coal. Consult your laboratory workbook. 




. . . Potassium and sodium are the 
names by which I have ventured to 
call the two new substances . . . 
. They agree with the metals in opaci- 
ty, lustre, malleability, conducting 
powers as to heat and electricity, and 
in their qualities of chemical combi- 
nation. Their low specific gravity 
does not appear a sufficient reason 
for making them a new class. 

Sir Humphry Davy, 1807 

Elements may be classified as metals or nonmetals. Thus far we 
have discussed a number of elements including oxygen, nitrogen, 
chlorine, bromine, iodine, fluorine, sulfur, and phosphorus. Each 
of these elements has more than one-half the number of electrons in 
its outer electron ring that are necessary to complete this ring. Hence, 
each is a borrower of electrons, has a negative valence, and has an ox- 
ide (or oxides) that is an anhydride of an acid. Each of these ele- 
ments is called a nonmetal. 

But the properties of these elements are not characteristic of all 
the elements. Certain elements have less than half the number of 
electrons in their outer electron rings necessary to complete these 
rings. Hence, each of these elements is a lender of electrons, has a 
positive valence, and possesses an oxide (or oxides) that is an an- 
hydride of a base. Each of these elements is called a metal. Of the 
100 chemical elements, 78 are metals. 

A comparison of sulfur and magnesium will help to make clear the 
differences between a metal and a nonmetal. 

This classification of elements into metals and nonmetals is very 
old, but it still continues to be of service. We must remember, how- 
ever, that certain elements behave either as a metal or as a nonmetal. 



a typical 


a typical 

a) 6 electrons in outside ring a) 2 electrons in outside ring 

b) borrower of 2 electrons b) lender of 2 electrons 

c) valence is 2 c) valence is +2 
Fig. 84. Structure of the atoms of a typical metal and nonmetal. 

How metals occur in nature. Metals occur either free (uncom- 
bined) , or chemically combined in compounds. Gold, silver, copper, 
and platinum are examples of metals that occur in the free state. 
The oxide of iron, the fluoride of aluminum, the chloride of sodium, 
the bromide of magnesium, the sulfate of barium, the carbonate of 
lead, and the phosphate of calcium are examples of compounds of 
metals that are found in the combined state. All these compounds 
are minerals, inorganic substances of definite composition found on 
the earth. A mineral such as mercuric sulftde, HgS, or rock salt, 
NaCl, from which an element, usually a metal, may be profitably 
extracted is an ore. 

How metals are extracted from ores. In mining operations, a 
mineral is first separated mechanically from the rock, or gangue 
(gang) , with which it is mixed by a process known as ore-dressing. 
The particular process used depends upon the differences in prop- 
erties between the ore and the gangue. Low-grade ores must be con- 
centrated to reduce the cost of extracting the metal. Various methods 
of ore-dressing, such as hand or gravity separation, leaching, mag- 
netic, and flotation, are described in the discussion of specific ores. 

The process of extracting a metal from its ore is called metallurgy. 
In general, the metallurgy of any metal depends upon the nature and 
purity of the ore, the properties of the metal to be extracted, and 
the cost of the processes. The four most widely used metallurgical 
processes are: (1) electrolysis; (2) reduction by means of carbon; 

(3) roasting, that is, heating the ore to change it to its oxide by 
oxidation (usually roasting is followed by reduction with carbon) ; 

(4) aluminothermy, that is, reduction with aluminum powder. 

The most widely used metal obtained by electrolysis is alumi- 
num. Iron is the chief metal obtained by reduction, and copper and 
zinc are the most common metals obtained "by roasting followed by 
reduction. Chromium and manganese are made by aluminothermy. 
These four metallurgical processes are discussed in Connection with 
aluminum, iron, copper, zinc, and chromium. 




Characteristic physical properties of a metal. Metals have many 
physical properties in common. A brief study of these characteristic 
properties helps us to realize why the classification of elements into 
metals and nonmetals has aided the development of chemistry. 

1) A metal is lustrous. That is, a metal has a definite shine. The 
luster of gold, silver, nickel, and copper is well known. 

2) A metal is malleable. That is, a metal can be hammered into 
thin sheets such as gold leaf and tin foil. The most malleable metal 
is gold, which has been beaten into sheets so thin that 300,000 of 
them placed one above the other make a pile only one inch thick. 
Carbon, a nonmetal, breaks when hammered; it is brittle. 

3) A metal is ductile. That is, a metal can be drawn into wire. 
Platinum, for example, can be drawn into such a fine wire that it 
cannot be seen by the unaided eye. The degree of ductility of metals 
varies greatly. Sulfur, a nonmetal, cannot be drawn into wire. 

4) A metal is a good conductor of heat. Place one end of a copper 
wire a foot long in a flame and notice how quickly your fingers hold- 
ing the other end of the wire become warm. 

5) A metal is a good conductor of electricity. Although silver is 
the best conductor of electricity, copper is used most generally be- 
cause of its much lower cost. On the basis of evidence now available, 
scientists think that good conductors of electricity, including metals 
in general and copper and silver in particular, are composed of 
atoms whose outer rings have only a few electrons loosely held. These 

(left) A prospector in Utah examines a sample of uranium ore. (right) Testing the 
ore with acid for the presence of copper and limestone. 

I'hotogmph*, Standard Oil Company (A r ./.) 

Pouring molten copper from an electric 
furnace into a mold. The "melt" con- 
sists of both virgin metal and high- 
grade scrap. 

/>'< xeurch Association 

electrons can, therefore, move along to the next atom, and so on, 
producing a flow of electricity. Nonmetals, such as sulfur, have in 
their outer rings a large number of electrons, not so free to move. 
Hence nonmetals are poor conductors of electricity. 

6) Other properties of a metal. All metals, with the exception of 
mercury, are solids at ordinary temperatures. They range in melting 
points from mercury ( 39 C.) to wolfram (3380C.) . Metals dif- 
fer widely in tensile strength, that is, the ease with which they can 
be pulled in two. Some steel has a tensile strength of 500,000 Ibs. 
per sq. in. cross section. 

They range in density from lithium (a little more than half as 
heavy as water) , whose density is 0.53, to osmium, whose density is 
22.5. Most metals are gray in color, the two most common excep- 
tions being copper (red) and gold (yellow) . They range in hard- 
ness from lithium, which is as soft as wax, to others which are very 
hard. All metals are crystalline. 

Carbide and Carbon Chemical* Company 

A flotation cell in which the ores 
of non-ferrous metals are con- 
centrated. Unwanted impurities 
settle to the bottom of the cell 
and the ore is skimmed off with 
the froth. 




Characteristic chemical properties of a metal. Metals also possess 
many chemical properties in common. 

1) Certain metals unite with oxygen, forming oxides that are basic 
anhydrides. The burning of magnesium, for example, produces mag- 
nesium oxide, which is the anhydride of magnesium hydroxide. 

2Mg -f O> - 2M gO; MgO + H 2 O -> Mg(OH) 2 

2) Certain metals unite with water, forming either bases or oxides, 
and liberating hydrogen. Thus steam passed over hot iron forms iron 
oxide arid hydrogen, whereas calcium reacts with water, forming 
calcium hydroxide and hydrogen. 

3Fe + 4H 2 O 
Ca + 2H 2 O 

> Fe 3 4 4- 4H 2 T 

> Ca(OH) 2 + H 2 

3) Certain metals decompose acids, liberating hydrogen or other 


Zn + 2HC1 

Cu + 2H 2 SO 4 

3Cu + 8HNO 3 

CuS0 4 + 2H 2 -f SO, t 

> 3Cu(NO s ) 2 4- 4H 2 -f 2NO 

4) Metals combine with nonmetals, forming salts. For example, 
they unite with chlorine, sulfur, and bromine, forming chlorides, 
sulfidcs, and bromides, respectively. 

5) Certain metals unite, with bases, liberating hydrogen. Thus, 
both aluminum and zinc react with sodium hydroxide, liberating 

2A1 + 6NaOH -* 2Na,<AlO 3 (sodium aluminate) -f 3H 2 T 
Zn -f 2NaOH -> Na 2 ZnO 2 (sodium zincate) + H 2 \ 

Copper and Brass Research Association 

These plates of impure 
copper are about to be sus- 
pended in a tank contain- 
ing copper sulfate and 
sulfuric acid. The copper 
will be refined by an elec- 
trolytic process described 
in Chapter 28. 


Metals may be listed in the order of their replacement power. 

It is generally known that gold does not tarnish in air and is not 
acted upon by any one acid, It is equally well known that iron, 
on exposure to air, is oxidized readily. We could, if it served any 
useful purpose, arrange all the metals in the order of their ability 
to resist oxidation. 

A more useful arrangement is based upon the ability of one 
metal to replace another from a solution of a salt of the latter. For 
example, when an iron nail is placed in a solution of copper sulfate, 
the iron becomes coated with a layer of pure copper, and iron sul- 
fate is formed. 

Fe + CuSO 4 - Cu + FeSO 4 

Similarly, copper placed in a solution of silver nitrate becomes 
coated with pure silver and will, itself, go into solution as copper 
nitrate. The reverse reactions will not take place under normal con- 
ditions. That is, a copper nail placed in an iron sulfate solution 
will not deposit iron. 

If we try these experiments, using a number of different metals 
and their salts, we can arrange the common metals in a definite 
replacement series. This replacement series of the common metals is 
also called the electrochemical series, and the electromotive series. 

How the electron theory explains the replacement series. Metals 
differ in their tendency to lose electrons and become ions. When 
free iron (Fe) replaces copper from a copper sulfate solution, the 
following electron reaction occurs: 

loses 2 electrons 

Fe + Cu ++ (SO 4 )~ -+ Fe++(SO 4 )-- + Cu 

s, ^ 

gains 2 electrons 

Electrically neutral iron loses two electrons and becomes ionic 
Fe + +, which is positively charged. Ionic copper from the copper sulfate 
solution gains two electrons, becomes electrically neutral, and precipi- 
tates out as free copper atoms. The sulfate ions remain unchanged. 



Zinc ^ 



OF THE | lead 











Fig. 85. A strip of metal, placed in a solution of a salt of a metal below it in 
the replacement series, replaces the less active metal, which precipitates onto the 

The only change that takes place, then, is a transfer of two electrons 
from free iron to ionic copper. Thus it seems that iron has a greater 
tendency than copper to lose electrons. That is, iron is more metallic 
than copper, and hence appears higher up in the replacement series. 
If we place iron in a solution of calcium chloride, no reaction takes 
place because the tendency of iron to lose electrons is less than that 
of calcium. 

Fe + Ca++ + 2C1~ -* no reaction 

The elements high up in the replacement series are so typically 
metallic and have such a great tendency to lose electrons, that even 
light causes them to throw off electrons. This fact is made use of in 
the photoelectric cell. Such a cell is frequently lined with a thin film 
of potassium, rubidium, or cesium, the most active metal known. 
When light strikes this film, it throws off electrons, which travel to 
a positively charged plate in the center of the cell. A very feeble 
electric current is produced. This feeble current, whose strength 
depends upon the intensity of the light that strikes the cell, may be 
amplified and thus made to control larger supplies of energy. This 
amplified current may close a switch, called a relay, which will start 
a motor and open a door, or count people going through a passage. 

The photoelectric cell is a vital part of sound-motion-picture and 
television equipment, and it is used also in the transmission of pic- 
tures by wire. Photoelectric cells are used in controlling the "blow" 
of Bessemer converters in making steel. 

Selenium, an element belonging to the sulfur family, was used in 
certain "electric eyes." Selenium is a good electric insulator in the 
dark, but in light it conducts an electric current to some extent. 
Later it was replaced by copper covered by a thin film of Cu,O. 

vacuated bulb 

cesium film 

Fig. 86. Construction 
of a photoelectric cell. 


The replacement series of the nonmetals. In studying the halo- 
gens, we learned that they, too, could be arranged according to 
their ability to replace one another from solutions of their salts. 
Thus, when chlorine is added to a solution of sodium bromide, 
bromine is liberated and sodium chloride is formed. According to 
the electron theory, this reaction is explained as follows: 

gains 1 electron 

SNa+Cl" + Br 2 

loses 1 electron 

Chlorine is more typically a nonmetal than is bromine. There- 
fore, it has a greater tendency to gain electrons. Free chlorine, which 
is neutral, takes one electron from the bromine ion and changes to 
Cl~, that is, it goes into solution. The bromine ion, on the other 
hand, after losing one electron, becomes electrically neutral, changes 
to the atomic form, joins with another atom of bromine, and is liber- 
ated as a free bromine molecule. Other nonmetals also may be 
grouped in a series according to their replacement powers. 

The replacement series of the metals is a useful tool. Under- 
standing the replacement series of the metals is of great value in 
studying these elements. Those above iron in the series, are so very 
active that they are never found free, while those below iron occur 
both in the free state and chemically combined. As we go down the 
series, the tendency of the metal to lose electrons diminishes, and 
hence the tendency to oxidize and to react with water or acids 
diminishes also. Thus gold and platinum, which are at the bottom 
of the list, do not oxidize in air even when hot, and are not attacked 
by water or even by any one acid. Therefore, knowing the position 
of a metal in this series, you can predict fairly well its chemical prop- 

Why hydrogen is included in the replacement series of the metals. 
Hydrogen, which from its physical properties could never be con- 
sidered a metal, belongs in the replacement series because its ion is 
usually positively charged and behaves chemically as a metal. All 
metals above hydrogen in the replacement series liberate hydrogen 
from acids. Those metals below hydrogen require oxidizing acids 
to dissolve them and liberate gases other than hydrogen from these 
oxidizing acids. For example, when copper reacts with nitric acid, 
nitrogen dioxide is given off. 

Humphry Davy isolates potassium and sodium. At the head of 
the replacement list of the metals are the alkali metals, lithium, 

Sir Humphry Davy (1778-1829), the son of 
a poor wood carver, was born at Penzance, 
Cornwall. The poet, Samuel Coleridge, de- 
clared that if Davy "had not been the first 
chemist, he would have been the first poet 
of his age." 

National Portrait Gallery, London 

potassium, and sodium. Two of these elements had been known for 
a long time as part of the alkaline compounds, potash and soda, be- 
fore they were isolated in a pure state. The alkali metals were orig- 
inally found in the ashes of certain plants. The name was taken from 
the Arabic al, meaning the, and quili, meaning ashes. 

Because of the extreme activity of these metals, many unsuccessful 
attempts had been made to isolate them. After the discovery, in 1800, 
of the galvanic current and the invention of the electric battery 
which soon followed, scientists used this new force in an effort to 
isolate sodium and potassium. 

Humphry Davy, an eminent English chemist, rose from humble 
beginnings to knighthood. He was the first to prove chlorine an 
element and, incidentally, he was versatile enough to spend his lei- 
sure hours writing fairly good poetry. In 1807 he sent the energy of 
150 electric cells through molten potassium hydroxide. At the nega- 
tive platinum electrode, Davy saw globules of a silvery substance 
form, and then spontaneously catch (ire. "His joy knew no bounds, 
he began to dance, and it was some time before he could control 
himself to continue his experiments." 

London received Davy's isolation of potassium as another wonder 
o the world, and he was lionized. Some people paid 100 dollars to 
attend his lectures on chemistry. Soon afterward, Davy obtained free 
sodium in the same way, and lithium from fused LiCl. 

How sodium is prepared for industrial use. The most recent 
method of obtaining sodium in large amounts differs somewhat from 
the method originally used by Davy. An electric current is sent 
through melted sodium chloride in a cell, such as is shown in Fig. 


sodium collect! 


to DC source 


melted NaCI 
metal screen 
MM head 

Fig. 87. Downs cell for 
the preparation of sodium. 
The cell was invented in 
1924 by J. C. Downs, an 
American chemist. 

iron or copper cathode 
>$* graphite anode 

87. Sodium ions travel to the iron or copper cathode, gain electrons, 
become sodium atoms, and collect as a mass of metallic sodium, 
which is drawn off from time to time. Chlorine ions travel to the 
graphite anode, lose electrons, and become gaseous chlorine, which 
leaves the apparatus as shown. This entire process is continuous. Po- 
tassium and lithium may be prepared in this same way using melted 
potassium chloride and lithium chloride. 

Physical properties of potassium and sodium. Sodium is a soft, 
silvery metal that melts just below the boiling point of water. Po- 
tassium, which is also soft and silvery white, melts at an even lower 
temperature. They are both lighter than water. Strangely enough, 
these two solids when mixed form a liquid alloy at ordinary tem- 

Some chemical properties of sodium and potassium. The electron 
pictures of lithium, sodium, and potassium are shown in Fig. 88. 

From these pictures we can tell that the valence of each of these 
elements is one. Each reacts with nonmetals, forming salts, and, on 
exposure to the oxygen of the air, each is quickly tarnished with a 
coating of its oxide. 

Since sodium has only one electron to lend, and oxygen must 
borrow two electrons to complete its outer ring, two sodium atoms 
combine with one atom of oxygen, and the oxide of sodium is, 
therefore, Na 2 O. Sodium peroxide, Na.X).,, (Na O O Na) , is 
formed when sodium is heated in air free from carbon dioxide. 
Because of its extreme activity, sodium cannot be kept exposed to 
air or under water. It is usually stored under kerosene, because kero- 
sene contains no oxygen. At high temperatures both sodium and 
lithium combine with hydrogen to form hydrides which react with 
water liberating hydrogen. 

LiH + H 2 O - LiOH + H 2 1 

Fig. 88. Structure of the atoms of the alkali metals. 

Lithium .'""'"^* 

Sodium x C 

Potassium x^- "~ 




What has been said about the chemical activity of sodium applies 
also to other alkali metals, namely, potassium, lithium, cesium, and 

How the alkali metals are used. In a recent year 150,000 tons of 
sodium metal were used in the manufacture of several compounds, 
such as sodium peroxide, sodium cyanide, sodamide (NaNH,,) used 
in making indigo, sodium hydride (NaH) used as a reducing agent 
in removing surface oxides from steel, and several detergents. A 
sodium-lead alloy is used in the manufacture of tetraethyl lead. 
Liquid sodium, because it is an excellent heat conductor, is em- 
ployed as a coolant in some nuclear reactors. Sodium is also used in 
the hot cathode sodium vapor lamp, which gives twice as much 
light as the common filament electric lamp using the same amount 
of current. This type of lamp is used chiefly in outdoor lighting. 

How can we test for the ions of sodium, potassium, and lithium? 
If a clean platinum wire is dipped into a salt of potassium and then 
placed in a nonluminous bunsen (lame, the flame becomes violet in 
color. The flame of all sodium salts is a distinct yellow; that of all 
lithium salts is red; cesium gives a bright blue flame. Since the pres- 
ence of even a trace of a sodium salt will obscure the violet color 
of potassium, the flame of a potassium salt can frequently be de- 
tected only when viewed through a piece of blue cobalt glass, which 
absorbs yellow light rays. 

(left) Removing a sodium brick from a shipping drum, (right) Making sodium 
pellets for laboratory use. Note the protective equipment in each case. 

Ethvl Corporation 

The bright line spectrum of lubricating oil (left) and of low-grade coal. 

Sucli flame tests are used on other metals besides the alkali metals. 
Thus, heated copper imparts a green color to the flame, and calcium 
gives the flame an orange-red coloration. 

How is a spectroscope used? In 1854, David Alter, a Pennsylvania 
physician, described a method of detecting an element by the color 
that it imparts to a flame. He also predicted the use of this method 
in determining the presence of elements in the sun. 

Five years later, Bunsen and Kirchhoff (kirK/hof) devised an in- 
strument called the spectroscope, which has since become a very 
powerful tool in the hands of chemists, physicists, and astronomers. 
In I860, with the aid of this instrument, two new elements, cesium 
and rubidium, were detected by Bunsen in a few grams of salt ob- 
tained by the evaporation of 40 tons of spring water. 

Nine years later, the element helium was discovered with the aid 
of the spectroscope by Janssen and Lockyer independently. Helium 
was found not on the earth, but in the sun, more than 90 million 
miles away. Before the close of the century, this new element was 
found on our own planet by Ramsay. 

With the aid of the spectroscope, other elements, present in such 
minute quantities that they had heretofore escaped discovery by 
even the most delicate instruments of science, were finally brought 
to light. Today, the spectroscope is used also in the study of the 
complex structure of the atom. 

This spectfOf rctpft makes Him records of spec- 
tra, simitar to shown bove From tuch 
tfMKtra, the composition of compounds may 

An analytical chemist using elec- 
tron diffraction equipment to 
identify crystalline material such 
as nickel oxide. Beams of elec- 
trons aimed at the sample break 
into a pattern which is made 
visible by the equipment. 

The principle of the spectroscope. To understand the principle 
of the spectroscope, we must understand a few things about light. 
Isaac Newton, in 1672, performed a classic experiment. He let a 
beam of sunlight pass through a narrow slit into a dark room and 
placed a glass prism in its path. A band of colors called the spectrum 
was formed. This can be explained by remembering that light is a 
form of energy which is transmitted by waves. 

Sunlight is made up of light of various colors. Each color has a 
different wave-length. Red has the longest wave-length (0.0000(i8 
cm.) , and violet has the shortest wave-length (0.000040 cm.) . The 
glass prism bent and split up, or refracted, the sunlight. The light 
which was refracted, or bent, least was the red, and that which was 
refracted most was the violet. Study the spectrum shown in the illus- 
tration in color following page 382. 

Though some self-luminous sources, such as the sun, have a con- 
tinuous spectrum, as shown in the illustration, an incandescent va- 
por or gas, such as heated sodium vapor or electrified neon gas, has 
a discontinuous, or bright-line) spectrum. The glowing vapor of each 
element has its own characteristic colored band of light. Thus, so- 
dium vapor has one bright yellow line, lithium has one red and one 
yellow line, and the vapor of iron has several hundred lines. 

The spectroscope, the most essential part of which is a glass prism, 
makes possible the quick analysis of incandescent vapors, and the 
detection of the smallest trace of an element. Less than a millionth 



of a milligram of sodium, and minute traces of poisons in blood 
can be detected by spectroscopic analysis. A spectroscope is thus a 
most useful tool to the chemists who specialize in the analysis of 
many kinds of substances. Such specialists are called analytical chem- 

Spectroscopic work of great precision is carried on today by means 
of the spectrograph. This instrument differs from an ordinary spec- 
troscope in that the observing telescope is replaced by a camera, 
which makes a photographic record of the spectrum under examina- 
tion. This permanent photographic record makes possible a more 
careful analysis of the spectrum. Spectographic analysis is used in 
steel and other alloys and compares favorably with the usual routine 
quantitative analysis. Analysis of matter by the study of spectra is 
called spectrum analysis. Because chemical manipulations are un- 
necessary, all measurements may be made quickly. 


Chemistry, Jan., 1945, pp. 37-43, and May, 1952. Published 
by Science Service, Washington, D.C. Contain the original 
papers describing the discovery of lithium, sodium, potassium, 
cesium, and rubidium, and more information on the other 

Mills, John. Electronics: Today and Tomorrow, pp. 84-94. 
D. Van Nostrand Co., New York, 1944. Tells the story of the 
photoelectric cell. 

Pough, Frederick H. A Field Guide to Rocks and Minerals. 
Houghton Mifflin Co., Boston, 1953. An excellent book for the 
boy or girl interested in mineral-collecting as a hobby. Attrac- 
tively illustrated. 


1. A metal is an element that lends electrons. It has a posi- 
tive valence, and its oxide is a basic anhydride. 

2. An ore is a mineral from which an element may be 
profitably extracted. 

3. Ore-dressing is the separation of a mineral from the 
valueless rock, or gangue, with which it is mixed. 

4. Metallurgy is the process of extracting a metal from its 
ore. The four chief metallurgical processes are (1) electrolysis, 
(2) reduction by means of carbon, (3) roasting, or heating, 


the ore to change it to its oxide, and (4) aluminothermy, 
using aluminum powder. 

5. Metals have certain characteristic physical properties. 
A metal is (1) lustrous, (2) malleable, (3) ductile, (4) a good 
conductor of heat, (5) a good conductor of electricity, (6) a 
solid mercury is an exception, (7) crystalline, and (8) vari- 
able tensile strength. 

6. Metals have certain characteristic chemical properties. 
Some (1) unite with oyxgen, forming oxides which are basic 
anhydrides, (2) unite with water, forming either bases or 
oxides with the liberation of hydrogen, (3) decompose acids, 
liberating either hydrogen or some other gas, (4) combine 
with nonmetals, forming salts, and (5) unite with bases, lib- 
erating hydrogen. 

7. Metals may be so arranged in a replacement series that 
each metal in the list will replace each metal below it from a 
solution of its salt. According to the electron theory, this be- 
havior is explained by the fact that the metals at the top of 
the list lose electrons more easily than those at the bottom, 
and hence go into solution more readily. The metals lower 
on the list take these lost electrons, become electrically neutral, 
and precipitate out as free metals. 

8. Some nonmetals, such as the halogens, may also be ar- 
ranged in a replacement series. According to the electron the- 
ory, chlorine has a greater tendency to borrow electrons than 
the other halogens and, hence, borrows electrons from those 
below it in the series. 

9. The flame test for identifying a metal consists of heating 
a metal or one of its salts in a flame and noticing the color that 
it imparts to the flame 

10. The spectroscope is an instrument devised by Bunsen 
and Kirchhoff. It is used to detect minute traces of elements 
in incandescent vapors. The most essential part of a spectro- 
scope is a glass prism, which disperses or breaks up the light 
into colored lines which are characteristic for each element. 

11. Spectrum analysis has been used in detecting the pres- 
ence of rare elements, such as the inert gases of the atmosphere. 
It is also used in the study of the structure of the atom. 


Group A 

1. According to the electron theory, how do metals differ 
from nonmetals? 


2. What are four typical nonmetals and four typical 

3. (a) Name a few metals that occur free, (b) Compare 
their chemical activity with the chemical activity of metals 
found combined. 

4. (a) In what way does a mineral differ from an ore? 
(b) Give an example of each, (c) Do minerals ever become 
ores? (d) Illustrate. 

5. (a) What are the four most general methods used in 
metallurgy? (b) Give an example of an element extracted by 
each of these methods. 

6. What are the characteristic physical properties of metals? 

7. (a) Arrange these eight metals in a replacement series: 
Zn, H, Cu, Na, K, Pt, Au, and Pb. (b) On your list, check 
those that will replace the hydrogen of dilute HC1. 

8. Arrange the number of each metal opposite the letter 
of the property of which it is an outstanding example. 

a) Best conductor of electricity 1) Cu 

b) Extremely ductile 2) Au 

c) Most malleable of all 3) Li 

d) Lightest metal 4) Hg 

e) Heaviest metal 5) W 

f) Liquid metal 6) Os 

g) Extremely high melting point 7) Ag 
h) Reddish luster 8) Pt 

9) Na 

9. (a) Write five equations illustrating five chemical prop- 
erties of metals, (b) In each case state the property. 

10. According to the electron theory, why does an iron nail 
become coated with Cu when placed in a solution of CuSO 4 ? 

11. In which of the following would a replacement reaction 
take place? Complete the equations for such replacements. 

a) Zn + Hg(NO 3 ) 2 -* e) Cu + AgNO 3 - 

b) Zn + CuSO 4 -> f) Cu + ZnSO 4 -> 

c) Zn + AgNO 3 - g) Cu + Hg(NO 3 ) 2 -> 

d) Zn -f Pb(NO 3 ) 2 - , h) Cu + Pb(NO 3 )a - 

12. Arrange a replacement sefies of some nonmetals you 
have studied. 

13. According to the electron theory, explain why Br 2 lib- 
erates free I 9 from a solution of KI. 


14. An element X is not found free. It attacks warm water, 
liberating H 2 , and tarnishes readily in air. Where would you 
place it in the replacement list of metals? See Tat>le 12. 

15. (a) What are two elements other than Na that belong 
to the sodium family? (b) Why are they called alkali metals? 

16. (a) How and by whom were Na and K first isolated? 
(b) How was the news of this achievement received? 

17. (a) What are three ways in which Na is similar to K? 

(b) one way in which it is different? 

18. By a diagram, show the present method of obtaining 
free Na. 

19. Make a diagram of the atom of Li, and, from this 
diagram, state its (a) valence, (b) chemical activity, and 

(c) atomic number. 

20. Using an equation, describe the action of Na on H 2 O. 

21. What weight of NaOH must be decomposed to produce 
69 g. of pure Na? 

22. A piece of Na is placed in H 2 O, and 336 ml. of H 2 are 
collected. What weight of Na took part in this chemical 

23. Which will require more H O in dissolving completely 
45 g. of Na or 79 g. of K? 

24. Copy and complete. Do not write in this book. Na is 
stored under .... When Na is exposed to air, the formula of 
the compound formed is .... Na is used as a catalyst in the 
preparation of .... An instrument that makes use of the ease 
with which Cs loses electrons is the .... The color imparted 
to a flame by K vapor is . . ., Li a ... color, and Na a ... 

25. (a) List four uses of sodium, and (b) three uses of 
other alkali metals. 

26. How is the spectroscope used in astronomy? 

27. Name three elements discovered by means of the spec- 

28. How can the spectrum of sunlight be obtained? 


29. What is the difference in the appearance of the spectrum 
of a luminous solid and the spectrum of an incandescent 

30. Hydrogen is included in the replacement series of the 
metals. Why? 

Group B 

31. Zn appears higher than Fe in the replacement series. 
What is the reason for coating Fe with Zn to prevent corrosion? 

32. Devise an experiment for obtaining Cu from CuSO 4 

33. Na and K kept under kerosene for some time lose their 
silvery luster. Explain. 

34. Discuss the use of spectroscopy in crime detection. 


1. Purchase or borrow a photoelectric cell and connect it 
in such a way that, when a flashlight is directed against it, a 
bell will ring. Explain its action to the class. What use would 
you make of a photoelectric cell? 

2. Make your own replacement series of some or all of the 
following metals: aluminum, copper, chromium, lead, tin, 
zinc, and calcium (if you can obtain a small piece from your 
teacher) . Use whatever suitable chemicals you can find around 
the house such as vinegar. Report your results. 

3. Prepare an illustrated ten-minute talk on mineral-col- 
lecting as a hobby. Bring some of your specimens to class. 

4. Write a 300-500 word essay on one of the various ana- 
lytical methods used by chemists today such as (a) wet 
method, (b) spectrophotometry, (c) chromatography, (d) tracer 




. . . / believe I speak for the vast ma- 
jority of all scientific men. Our ob- 
ject is riot to make fobs and divi- 
dends. These are a means to an end, 
merely incidental. We wish to abol- 
ish drudgery, discomfort, and want 
/row the lives of men, and bring 
them pleasure, comfort, leisure and 
beauty. Harold C. Urey, 1934 

The world receives a valuable gift. In 1825, Hans Christian 
Oersted (lir'stetfe) , a Danish scientist, announced that he had isolated 
aluminum by gently heating aluminum chloride and potassium 
amalgam. In 1827 Woehler repeated Oersted's experiments without 
success. Woehler finally obtained aluminum by substituting metallic 
potassium for the potassium amalgam used by Oersted. 

One of the first of Woehler's American students in Germany was 
Professor Jewett, of Oberlin College. He brought back to America 
the story of Woehler's isolation of that extremely light, silvery metal, 
aluminum, fewett was fond of talking to his classes of this strange 
metal, which no one had as yet been able to obtain inexpensively in 
spite of its great abundance in minerals. 

One day, as Professor Jewett spoke of the fortune that awaited the 
man who was able to develop a simple method for extracting alu- 
minum, one of the students nudged his young classmate, Charles 
Martin Hall. Chemistry had captivated Hall, and his classmates 
had known him to make all sorts of experiments, hoping to make a 
great discovery some day. Here was his chance. His response to that 
nudge was, "I am going after that metal," and Hall went to work at 
once in his father's woodshed 



Hall attacked his problem scientifically. He knew that only the 
most active metals, such as sodium and potassium, were reducing 
agents strong enough to liberate aluminum from aluminum chloride. 

A1C1 3 + 3K -> 3KC1 + Al 

Potassium had been used in the method developed by Oersted 
arid used later by Woehler, and sodium had been substituted tor po- 
tassium by the French chemist, Henri Sainte-Claire Deville (saNt- 
klar' de-vel') . But both potassium and sodium were too expensive 
to use in a commercial method. Hall finally discarded all methods 
that depended upon the action of a reducing agent and turned to 
electrolysis, in spite of his knowledge that Davy, who had isolated 
the alkali metals by electrolysis, had failed to get pure aluminum in 
this way. 

Aluminum oxide, called alumina, was the natural starting point. 
Alumina in hydrated form is the chief component of bauxite, the 
richest ore of aluminum. But alumina has an extremely high melting 
point. To melt alumina was commercially impracticable. But if an 
electric current was to liberate free aluminum from it, alumina had 
to be either melted or dissolved. Perhaps (the thought came to Hall 
in one of those flashes of genius) some mineral that would act as a 
solvent for aluminum oxide might be found. After trying a number 
of minerals, he came across a milky-white, glassy solid called cryolite. 
He melted this with some difficulty and then threw in some alumina. 
The alumina dissolved readily. He passed a current through the 
solution of alumina in cryolite and, to his intense joy, found that 
metallic aluminum was deposited at the cathode 

On February 23, 1886, he burst excitedly into the laboratory of 
Professor Jewett and, holding a few aluminum "buttons," exclaimed, 


Charles Martin Hall, the discoverer of the 
electrolytic process of producing alumi- 

p, o c ,- , + I. .1. 4. ^busbar 

Fig. 89. Cross section of a '.. \ m \ i,i D ..' ^,^ Mt JL A! o 

* , .. . . . . M TT T rTx^'^crusF or AUw 3 

furnace in which aluminum is I] 1 _ | _ I ^Jr[ 

produced commercially by the ^j^jjjAjjjLj^^ AI 2 O 3 In 

Hall process. carbon anode + ^g^^^^^J^Tn fused^cryolite 

carbon lining cathode aluminum 

"I've got it!" Hall soon obtained a patent on his process and two 
years later the Pittsburgh Reduction Company, which later became 
the Aluminum Company of America, was founded. In 1914 Hall 
died, world-famous and a multimillionaire. He left most of his for- 
tune to Oberlin College and other educational institutions. 

Discoveries often result from social needs. Hall was 22 when he 
produced aluminum. Exactly two months after Hall had handed his 
teacher the first samples of his aluminum, another chemist, Paul 
Heroult (a-roolt') , also 22 years of age, applied for a patent in France 
on an identical commercial method of preparing aluminum. 

This is not a rare example of simultaneous discoveries in the his- 
tory of science. Advances in science often are made in different parts 
of the world at almost the same time. They are frequently the final 
step in a long series of experiments conducted by many research 
workers in widely separated laboratories. The scientist who is fortu- 
nate enough to publish his discovery first is recognized as the hero 
of a battle in which many other soldiers have been engaged. The 
heroes of science, on the whole, concede this element of good for- 
tune. Can you think of another such instance in the story of scientific 

Metallurgy of aluminum by the Hall process. Hall's process is still 
in use. The electric furnace used is an open cell about 25 to 30 feet 
long, lined with carbon, which constitutes the cathode. Powdered 
cryolite mined in Greenland or made synthetically is placed in the 
cell, and as the current passes through it, it melts. Aluminium oxide, 
or alumina, is a white powder obtained by refining bauxite ore. It 
is added to the molten cryolite and immediately dissolves. The 
aluminum oxide dissociates, forming positive aluminum ions and 
negative oxygen ions. 

A1 2 O 3 - 2A1+++ + 3O 

Carbon rods are suspended in the molten aluminum oxide solution 
and act as the anode. When the circuit is closed, the aluminum ions 
travel to the cathode, where they obtain electrons which change 
them from aluminum ions to free aluminum. This free molten alu- 
minum then settles to the bottom of the cell. Later, a hole at the 



bottom of the cell is unplugged, and the molten aluminum is tapped 
off into large ladles and cast in molds, in which it solidifies as pig 
aluminum. The oxygen ions, in the meantime, have traveled to the 
anode, given up their electrons, and changed to free oxygen. This 
oxygen combines with the carbon of the anode and forms carbon 
dioxide, given off as a gas. 

c -f o 2 - co 2 T 

The process is continuous. Aluminum oxide is added, aluminum is 
removed, and the carbon anodes are replaced from time to time. 
The original cryolite, Na 3 AlF, though it contains aluminum, does 
not decompose. It acts only as a solvent. Many of these electrolytic 
cells are joined in series. Commercial cells produce about 500 pounds 
of 99+ percent pure aluminum per day. 

Bauxite, found in large amounts in Surinam (Dutch Guiana) , 
and British Guiana Arkansas leads the United States in the pro- 
duction of this ore contains a fair percentage of the oxides of iron, 
silicon, and titanium. If these impurities are not removed before 
the bauxite is added in the electric furnace, the aluminum produced 
is impure. 

After the aluminum is drawn from the electrolytic cells, the pig 
metal is remelted so that the nonmetallic impurities may be skimmed 
off. If aluminum alloys, rather than pure aluminum, are desired, the 
alloying may be done during the remelting. The chief alloying ele- 
ments include copper, magnesium, manganese, silicon, zinc, iron, 
nickel, and chromium. 

The physical properties of aluminum. Aluminum is silvery white 
in color and is one of the lightest of the common metals. It is only 
one-third as heavy as iron. It is very malleable and ductile and com- 
pares well with both silver and copper in the ease with which it con- 
ducts both heat and electricity. It can be worked readily; that is, it 
can be cast, rolled, forged, extruded, machined, or drawn. Parts can 
be joined by welding, brazing, and riveting. 

The chemical properties of aluminum. The atomic weight of alu- 
mium is 27. Its atomic number is 13; hence it has only three elec- 
trons in its outside ring. It is, therefore, a metal with a valence of 
plus three. Aluminum is an amphoteric element and may act as either 
an acid or a base. It is attacked by strong bases as follows: 

2A1 + 2NaOH + 2H 2 O - 3H 2 1 + 2NaAlO 2 (sodium aluminate) 

This sodium aluminate is the salt of aluminic acid, H 3 A1O 3 . Be- 
cause of the reaction between aluminum and the strong bases, or 



substances with basic reactions, such as washing soda, care should be 
taken not to heat such substances in aluminumware. 
Aluminum is attacked by nearly all acids, forming aluminum salts. 

2A1 + 6HC1 - 3H 2 1 + 2A1C1 3 

The surface of aluminum oxidizes rapidly in air, forming alu- 
minum oxide, AL,O 3 . This extremely thin, transparent, but tough 
film acts as an excellent protective coating and, unlike iron rust, 
adheres firmly to the surface of the metal, thus preventing further 
oxidation unless the coating is perforated. Alclad is a sheet of alu- 
minum alloy such as duralumin covered with a layer of aluminum. 
It resists corrosion very well. 

Tremendous growth of the aluminum industry. Before the Hall 
process was introduced, aluminum was not used widely, because of 
the great cost of preparing it. It is a far cry from the world produc- 
tion of two tons of aluminum in 1859 at 17 dollars per pound to the 
more than one million tons of this metal produced in the United 
States and Canada alone in a recent year at about 20 cents a pound. 

The Hall process gives primary aluminum, that is, metal pro- 
duced directly from an ore or ores. But a significant and perhaps 
increasing source of aluminum, and other metals as well, lies in 
secondary sources,, that is, sources from which a metal is recoverable 
from one use for reuse in another. Secondary or scrap aluminum is 
a very large source of pure aluminum. Production of metals from 

Surface mining of bauxite in Surinam. Over-lying earth is first removed. The ex- 
posed ore is then loosened by blasting and loaded into the mine cars. 

Aluminum Company of America 

An "aluminum skyscraper" in Pitts- 
burgh, Pa. This 30-story building is 
constructed of aluminum panels 
mounted on a steel framework. The 
ceilings, wiring, ventilation ducts, 
doors, hardware and most of the 
plumbing are made of aluminum. 

secondary sources is one phase of an intelligent metals conservation 
program. Such a program can do much to conserve natural resources. 

By far the largest users of aluminum are the transportation indus- 
tries. Great quantities of the "metal with wings" are used in the 
construction of airplanes, streetcars, railroad cars, locomotives, steam- 
ships, motorships, automobiles, trucks, buses, bicycles, and motor- 

Through decreasing the weight of such carriers as airplanes, rail- 
road cars, and trucks, payload and, thereby, revenue can be increased. 
At the same time, if the reduction in weight applies only to the total 
gross weight of the vehicle, as in an automobile or bus, much less 
energy is required to attain and maintain speed, much less energy 
is lost in stopping, and operating costs are thereby reduced. 

Electricity and aluminum. The electric industries use thousands 
of tons of aluminum yearly in lines for the transmission of electricity 
over long distances. For this purpose, aluminum cable with a steel 
reinforcing core is used instead of copper. Since aluminum is lighter 
than copper, fewer towers are required to support the cables. More 
than 1.5 million miles of aluminum transmission cables carry elec- 
tricity to almost all parts of this nation. Aluminum is used also in the 
production of parts for electric equipment of many kinds, such as 



vacuum cleaners and various household appliances, and particu- 
larly in parts for radios and other electronic equipment. 

Aluminum in the kitchen. One use of aluminum goes back to 
1890, when the first aluminum cooking utensils were produced. 
Since aluminum is an excellent conductor of heat, and at the same 
time is very light, aluminum cooking utensils are very popular. 

Certain alkaline foods and waters heated in aluminum may pro- 
duce a superficial discoloration, which is harmless and readily re- 
moved by a mild abrasive cleaner. Do not permit all the cooking 
liquid in a lightweight aluminum utensil to boil off, for if this 
occurs, a hole may be "burned" in the bottom of the utensil as a 
result of the relatively low melting point of aluminum. 

Because strong alkalies attack the protective coating of alumi- 
num oxide that forms on the surface of aluminum as well as the 
aluminum itself, cooking utensils made of the metal should not be 
scrubbed or polished with harsh alkali cleaners. To clean aluminum 
utensils use soap and water or mild abrasive cleansers only. 

Other uses of aluminum. In the packaging of foods and other com- 
modities, aluminum foil has almost entirely replaced tin foil. Candy 
bars, chewing gum, cream cheese, camera film, and countless other 
articles go to market in shining dress. Aluminum foil coated with a 
plastic film is suitable for the packaging of almost any kind of food. 
Aluminum leaf is used in photoflash lamps. 

Collapsible tubes made of aluminum carry such items as shaving 
cream, tooth paste, and cosmetics, while vital serums and various 
other pharmaceutical preparations are packed in glass bottles with 
aluminum seals. Aluminum paints are widely used for protecting 
both wood and metal. Aluminum foil is used for home insulation. 

Aluminum Company of America 

Circular aluminum blanks 
being removed from the 
conveyor of an annealing 
furnace. These blanks will be 
used in the manufacture of 
cooking utensils. 





Much furniture and many decorative articles for the household are 
made of aluminum. 

The buildings of Rockefeller Center in New York contain more 
than 1000 tons of aluminum in the vertical panels between windows. 
In the finishing of steel, large quantities of aluminum are used in 
removing oxides from the molten steel. The largest use of aluminum 
is in the form of alloys of much greater strength than pure aluminum. 

Thermit is used in welding. Attempts to reduce aluminum oxide 
with carbon failed because of the great attraction of aluminum for 
oxygen. Aluminum is a powerful reducing agent, especially when 
it is in the form of a fine powder. Because a powder has a much 
greater reacting surface than a solid, a powder makes possible a more 
intense chemical reaction than the same weight of the same solid in 
larger pieces. 

When a mixture of powdered aluminum and iron oxide is ignited 
by means of a fuse, such as a strip of magnesium ribbon, a chemical 
reaction takes place at once in which the aluminum takes the oxygen 
away from the iron oxide, leaving a residue of pure iron. 

8A1 + 3Fe 3 O 4 - 4A1 2 O 3 + 9Fe 

The heat of this reaction is so great that the iron formed is molten. 
This mixture of aluminum and iron oxide, known as thermit, is used 
in welding broken propeller shafts, rudder frames, locomotive parts, 
and in situations where repairs must be made on the spot. It is also 
used in one type of incendiary bomb. 

The metals chromium and manganese (and wolfram, vanadium, 
molybdenum, silicon and boron) may be extracted from their oxides 
or ores by aluminothermy, an aluminum reduction similar to the 
thermit reaction. The equations for the reduction of chromium and 
manganese ore are: 

Cr 2 O 3 + 2A1 -> A1 2 O 3 + 2Cr 
3Mn 3 O 4 + 8A1 - 4A1 2 O 3 + 9Mn 

This is a common way of manufacturing or producing such metals. 
What are alums? When potassium sulfate is dissolved in a solution 
of aluminum sulfate, the two salts combine and crystallize out as 
potassium aluminum sulfate. 

K 2 SO 4 + A1 2 (SO 4 ) 3 + 24H 2 O -> 2KA1(SO 4 ) 2 12H 2 O 

This compound is called common alum. It is one of a group of salts, 
called the alums, which resemble one another in the eight-sided form 
of their crystals, in their solubility in water, and in their type for- 
mula. The type formula of an alum is XY (SO 4 ) L > 12H 2 O. In it X may 
be K, Na, or NH 4 , and Y may be a trivalent element, such as Al, Fe, 
or Cr. An alum is a double salt, that is, a salt containing two metals 
and one acid radical. 

Common alum has a sweetish taste and is used in "Foamite," mor- 
danting (see page 605) , and water purification. Alum is used also 
in making alum baking powder and as an astringent, a substance that 
contracts skin tissues. It is used in the sizing of paper. Chrome alum 
is used in the tanning of leather, during which skins or hides are 
made softer, and more resistant to the action of bacteria. 

Aluminum hydroxide, its preparation and properties. When any 
soluble aluminum salt is added to water or to ammonium hydroxide, 
aluminum hydroxide, a white, gelatinous precipitate, may be formed. 
The aluminum salt is said to hydrolyze, since, as in the case of the 
solution of aluminum chloride in water, an acid and a base are 

A1 2 (SO 4 ) 3 + 6NH 4 OH 
A1C1 3 + 3H 2 O 

- 2A1(OH) 3 1 + 3(NH 4 ) 2 SO 4 

3HC1 + A1(OH) 8 1 


Use by approximate 

Industrial, agricultural, 
and mining machinery 


Aluminum hydroxide may act either as an acid or as a base. When 
it comes in contact with a base, it acts as a weak acid and combines 
with the base. 

HaAlOs + NaOH -> NaAlO 2 (sodium aluminate) + 2H 2 O 

A1C1 8 is the active constituent, about 15 percent, of many body 
deodorants and antiperspirants. 

Aluminum sulfate in the purification of water. As you know, 
aluminum sulfate, when it is added to water, hydrolyzes, forming 
jelly-like aluminum hydroxide. 

A1 2 (S0 4 ) 3 + 6H 2 - 2A1(OH) 3 j + 3H 2 SO 4 

This aluminum hydroxide gradually settles to the bottom and carries 
down with it any particles that are floating in the water, including 
bacteria, industrial wastes, and fine clay. This process, used in the 
purification of water, is called coagulation and is the first step in 
clearing water of its turbidity. Coagulation does not remove dis- 
solved impurities. 

Ferrous sulfate is sometimes used instead of aluminum sulfate, 
because it also forms a gelatinous precipitate, ferrous hydroxide, in 
water. But ferrous hydroxide rapidly oxidizes to ferric hydroxide, 
which is also gelatinous. Hence, it is the ferric hydroxide that actu- 
ally reduces the turbidity, or clarifies the water. 

FeS0 4 + 2H 2 -> H 2 SO 4 + Fe(OH) 2 j (ferrous hydroxide) 

Aluminum oxide ore, gem, abrasive, refractory. When alumi- 
num hydroxide is heated, it forms a white, insoluble oxide of alumi- 
num, which melts above 3600 F. 

2A1(OH) 3 -> A1 2 O 3 + 3H 2 O 

As hydrated oxides, A1 2 O 3 3H 2 O and A1 2 O 8 H 2 O are found widely 
distributed in the aluminum ore, bauxite. The precious gem stones 
ruby and sapphire are composed of alumina, colored by the pres- 
ence of small amounts of metal oxides. Successful methods have 
been developed to prepare synthetic rubies and sapphires by melting 
pure aluminum oxide in the heat of an oxy hydrogen flame. These 
synthetic stones cannot easily be distinguished from natural gems. 

The production of synthetic rubies and sapphires is increasingly 
important, for they are used as bearings in watches, electric indica- 
tors, sensitive electric relays, and in thousands of other kinds of pre- 
cision instruments. 


Emery is a natural aluminum oxide, which is extremely hard and 
can be used as an abrasive for grinding, polishing, drilling, and cut- 
ting. It is almost as hard as diamond. Fused alumina, for use as an 
abrasive, is prepared in large quantities by the fusion of alumina in 
an electric furnace. One trademark for such fused alumina is "Alun- 

Not only is fused alumina a good abrasive, but because of its high 
melting point, it is an excellent refractory. A refractory is a sub- 
stance which, because it melts at a high temperature and is chem- 
ically inert, can be used for furnace linings and in similar installa- 
tions which must withstand extreme heat. Fused alumina refractories 
are therefore used in making bricks, spark plugs, crucibles, cements 
for high-temperature work, and high-temperature thermometers 
called pyrometers. Because alumina is quite inert chemically, it is 
employed as a catalyst support, and for some laboratory ware such 
as porous plates used in filtering chemical solutions. 

Activated alumina is a highly porous A1 2 O 3 which is used to adsorb 
moisture in air conditioning, and to dry such gases as propane and 

Newcomers, the light alloys of aluminum. The American Society 
for Testing Materials defines an alloy as: A substance consisting of 
two or more metallic elements or of metallic and nonmetallic ele- 
ments which are soluble in each other when molten, and which do 
not separate into distinct layers when solid. The art of making alloys 
is very old. One of the oldest alloys is bronze. This alloy gave its 
name to an era of human progress, the Bronze Age, which began 
about 2500 B.C., and gave way to the Iron Age some 15 to 20 cen- 
turies later. 

Alloys are in many cases neither true solutions nor compounds. 
Whereas, for example, Cu and Au mix in all proportions, others 
have a limited solubility. Some alloys consist of metals present in 
almost constant proportions such as NaZn xl . The properties of an 
alloy may differ radically from those of the elements that compose 
it. An alloy is usually harder and more resistant than any of its con- 
stituents. For these reasons, alloys are extremely important. By mix- 
ing two or more metals, we may obtain an alloy whose properties 
are immensely more valuable than those of any of its ingredients. 
For example, the tensile strength of many aluminum alloys is greater 
than that of pure aluminum. 

Some of the alloys of aluminum containing copper, manganese, 
magnesium, or other metals are light, strong, and easily machined. 
Such alloys make ideal materials for the framework of the wings and 





fuselages, and landing gear and propellers of airplanes, and for cer- 
tain parts of automobiles and railway cars. "Duralumin," was one of 
the first strong aluminum alloys to be given additional strength by 
heat treatment. Other alloys, stronger, tougher, and more durable, 
have largely replaced it. 

Some of the modern alloys of aluminum are actually stronger than 
structural steel, and much lighter. Manufacturers of aluminum alloys 
call their products by many different names. Aluminum alloys of the 
Aluminum Company of America are called "Alcoa" alloys, and usu- 
ally a number indicates the particular "Alcoa" alloy. 

Whereas there are about 40 metals in common use, more than 
6000 alloys have helped make possible our modern industrial world. 
If we try to escape from alloys, we escape from civilization. Steel, 
the most important single alloy of all, is discussed in Chapter 27. 

Magnesium, lightweight of the structural metals. Magnesium is 
one of the metals most commonly alloyed with aluminum. It can also 
be used for many of the same purposes as aluminum with which it 
has become competitive. 

World War II, with its tremendous military demand for light met- 
als, saw the infant magnesium industry grow from a healthy baby to 
a full-grown adult almost overnight. Production of magnesium 
jumped in the five years following 1939 from 3500 tons to a high of 
nearly 200,000 tons. Volume for volume, magnesium in 1953 at 26.5 
cents a pound was less expensive than aluminum selling at 20.9 cents 
a pound. 

Magnesium is a soft, silvery-white metal about 60 percent as heavy 
as aluminum. In pure form it has little structural strength, but 
when properly alloyed, the resulting materials have good structural 
strength. Although the strengths of magnesium alloys are less than 
those of the heavier structural metals, by using slightly greater thick- 
nesses of the alloys, structural shapes as strong as steel, but with only 
a fraction of the weight of steel, may be made. 

"Dowmetal" is the trademark of a family of magnesium alloys 
made by the Dow Chemical Company, pioneer producers of magne- 
sium. All "Dowmetals" are composed of magnesium and varying 
percentages of other metals such as manganese, aluminum, or zinc. 
The percentages of these metals in the alloys vary with the char- 
acteristics desired in the alloy. For example, "Dowmetal A" contains 
eight percent aluminum and 92 percent magnesium. "Dowmetal G" 
contains ten percent aluminum, 0.1 percent manganese, and 89.9 
percent magnesium. A similar series of magnesium alloys is manu- 
factured by the American Magnesium Corporation. Such alloys are 
designated by number, for example, "AM3S" or "AM57S." Magne- 
lite, used in kitchenware, contains Al, Mg, and silicon. 

Chemical properties of magnesium. Magnesium is high in the 
replacement series of metals, but pure magnesium and some of its 
alloys corrode relatively slowly even in moist air. Like aluminum, 
magnesium is a self-protecting metal. The protective film is probably 
a basic carbonate. Powdered magnesium or thin strips of magnesium 
burn in air when ignited, forming magnesium oxide. Hot magne- 
sium burns in CO 2 to form MgO and C. 

2Mg + C0 2 -> 2MgO + C 

Fused magnesium oxide under the name of "magnorite" is a refrac- 

Magnesium from the sea. Magnesium is one of the most abundant 
industrial metals in the earth's crust. There is an almost unlimited 
supply of it in the sea, for one cubic mile of sea water contains nine 
billion pounds of this metal in the form of MgCl 2 . The brine wells, 
such as those near Midland, Michigan, are another source of magne- 
sium chloride. Until rather recently, magnesium production in the 
United States was limited almost completely to the electrolysis of 
molten magnesium chloride obtained from the sea or salt wells, or 
of fused carnallite, MgCL KC1. During electrolysis, the magnesium 

These workmen are pouring 
molten magnesium into a 
mold. Magnesium can be 
cast, molded, extruded, and 
worked by many of the com- 
mon methods of the metals 

collects at the cathode, and chlorine at the anode. Magnesium is thus 
the first structural metal obtained from the sea. 

Another source of magnesium is dolomite, MgCO ; , CaCO,, which 
is very widely distributed throughout the world, forming entire 
mountain ranges. Magnesium has been obtained from this ore by 
roasting it to form MgO, and then reducing the magnesium oxide 
at high temperatures with carbon. 


Holmes, Harry N. Out of the Test Tube. Emerson Books, 
New York, 1945. "Three Light Metals" are discussed in this 
very popular science book. 

Forbes, R. J. Man, the Maker. Harry Schuman, New York, 
1950. An excellent and compact history of technology and 
engineering including a short account of the Hall process. 

Wade, Frank B. "Man-Made Gems." Journal of Chemical 
Education, June, 1931, pp. 1015-1026. Describes the manufac- 
ture of artificial rubies and sapphires, and the methods used in 
detecting differences between natural and artificial gems. 


1. An alloy is a substance consisting of two or more metallic 
elements or of metallic and nonmetallic elements that are 
soluble in each other when molten and do not separate into 
distinct layers when solid. 



2. A double salt is a salt that contains two metals and one 
acid radical. 

3. Frequently, scientific discoveries and inventions are made 
independently and simultaneously in different laboratories. 


Group A 

1. By whom was Al isolated in 1825? 

2. (a) When and by whom was the first successful com- 
mercial method of producing Al devised? (b) What were the 
circumstances surrounding his solution of this problem? 

3. Name five elements that are prepared by electrolysis. 

4. Write an equation showing how Woehler obtained 
pure Al. 

5. In the metallurgy of Fe, C is used to reduce Fe 2 O 3 . Why 
is C not used to liberate Al from its oxide? 

6. (a) Make a labeled diagram of the electrolytic cell in 
which Al is freed from its oxide, (b) Discuss in detail the 
chemistry of the extraction of Al by the Hall process. 

7. (a) What is the function of Na 3 AlF in the Hall process? 
(b) What evidence is there that the Na 3 AlF 6 does not supply 
some of the Al formed at the cathode? 

8. (a) Is bauxite just as it comes from the mines used in 
extracting Al by the Hall process? (b) Explain. 

9. Give reasons why each of the following statements is 
either true or false: (a) O 2 is liberated at the anode of the 
aluminum cell, (b) The carbon anodes of the aluminum cell 
must be replaced from time to time, (c) Much Al is produced 
from alumina in plants located near Niagara Falls. 

10. Find the percentage of Al in (a) pure alumina and 
(b) in pure Na 3 AlF . 

11. Aluminum hydroxide may be written Al (OH) 3 or 
H 3 A1O 3 . Explain. 

12. Lustrous Al becomes covered with a coating in air. Ex- 
plain this chemical change. 

13. Al is higher in the replacement series of the metals than 
Fe, but it corrodes much less than Fe. Explain. 


14. Write the equations for: (a) Action of hydrochloric 
acid on aluminum, (b) Action of a strong base on aluminum, 
(c) What precautions are necessary in using aluminumware 
because of this action of bases? 

15. 300 Ib. of powdered Fe 3 O 4 are available. How much 
thermit mixture can be prepared from this Fe 3 O 4 ? 

16. (a) What properties of Al make it a satisfactory material 
for use in construction work? (b) Why cannot Al displace 
steel completely in construction work? 

17. (a) What are the most important uses of Al? (b) Op- 
posite each use write the properties that make this use pos- 

_ t . . . 

18. Make a list of ten articles used by you every day which 
contain Al. 

19. (a) What is an alloy? (b) Give three examples. 

20. What are the composition and uses of one of the chief 
alloys of Al? i 

21. How is a photoflash bulb constructed? 

22. (a) Of what is thermit composed? (b) For what is it 
used? (c) Write the equation for the reaction that takes place 
when thermit is ignited, (d) What is the source of the heat 
produced? (The temperature is so high that Fe, which melts 
at 1530C., becomes liquid.) 

23. (a) What two metals are extracted from their ores by 
the use of Al? (b) Explain, using an equation. 

24. (a) What is a double salt? (b) Give one illustration. 

25. Explain how it is possible to have an alum that does not 
contain Al. 

26. Write an equation showing how aluminum hydroxide 
is prepared in the laboratory. 

27. (a) What property of Al (OH) 3 makes it useful in puri- 
fying water? (b) Of what kind of impurities does it free the 
water? (c) How? (d) What is added to the water to form 
the Al (OH) 3? 

28. Explain how artificial rubies are prepared. 

29. (a) How is fused alumina made? (b) What is activated 
alumina? (c) What are some of its uses? 

30. (a) Describe two chemical properties of Mg. (b) For 
what is Mg used? 

31. (a) Wfcat metal is the chief competitor of Al? (b) Why? 


Group B 

32. (a) What is the effect of A1 2 (SO 4 ) 3 on moist blue litmus 
paper? (b) Explain. 

33. Describe one industrial process for the production of Mg. 

34. (a) What is a refractory? (b) Name two refractories 
that are used commercially in large quantities. 

35. Mg is a very abundant metal. Why has it only lately 
come into widespread use? 

36. One cu. ft. of pure water weighs 62.5 Ib. What does one 
cu. ft. of pure Al weigh? 

37. Washing soda should not be boiled in aluminum uten- 
sils. Why? 

38. Name three chemical discoveries that were made by 
young men. Do not include Hall's discovery. 

39. Compare the properties of an alloy with the properties 
of (a) a compound, (b) a solution, and (c) a mixture. 

40. (a) What kind of salts hydrolyze in H 2 O? (b) Explain. 

41. Al is the most abundant metal in the earth's crust. How- 
ever, Fe costs less than Al. Explain. 

42. (a) Write the reaction that occurs when A1 2 (SO 4 ) 3 is 
placed in water, (b) Write the reaction that occurs when 
Na 2 CO 8 is placed in water, (c) What ions exist in each of these 
solutions? Why? (d) If we add A1 2 (SO 4 ) s to a solution of 
Na 2 CO 3 , we get Al (OH) 3 and not A1 2 (CO 3 ) 3 . Why? 

43. The thermit process has been called a vestpocket blast 
furnace. Explain this comparison. 


1. Make a small model of the electrolytic cell used in mak- 
ing pure aluminum. 

2. Make a collection of articles or samples of alloys contain- 
ing aluminum. Report to the class on the composition and use 
of each item collected. 

3. At the hundredth anniversary of the opening of the U.S. 
Patent Office (1936) the 12 greatest inventions made in this 
country were listed. Low-cost aluminum was included. Make a 
list of 11 other inventions which you might have included in 
such a list. Consult your teacher of history. 




To an extent not generally appreci- 
ated, U.S. industrial strength is based 
on an accident of nature: the unique 
geographical combination of Minne- 
sota and Michigan ore, Appalachian 
coking coal and the Great Lakes 
highway. On this triple gift of nature 
rests our towering steel industry. To 
it we owe, in the final analysis, our 
standard of living. Leonard Engel in 
Scientific American, May, 1948 

Importance of steel in our machine age. In a recent year, more 
than 100 million tons of steel were produced to satisfy the appetite 
of the machine age. The iron produced in an entire year a hundred 
years ago would meet our present needs for just one day. This 
fact is easy to understand when we think of the great skyscrapers, 
the thousands of miles of rails, the millions of automobiles, the 
thousands of locomotives, ships, and bridges, and the great num- 
bers of machines made from iron which man did not have even 
a century ago. Iron ranks first among the metals in tonnage used 
and, next to aluminum, it is the most abundant metal in the earth's 

Where does this iron come from? Iron is not found free, except 
in meteorites that have fallen on the earth as visitors from outer 
space. Iron ores, however, are numerous and widely distributed. 
The most important iron ore deposits (chiefly hematite, composed 
largely of Fe 2 O 3 ) occur in the "ranges" near Lake Superior. The 
discovery of these huge deposits of iron had a greater effect on Ameri- 
can Iffe than the more romantic discovery of gold in 1849. One of 
the iron mines in the Mesabi range yields every two weeks a volume 
of iron equal to the Great Pyramid. Here the ore, containing about 




50-55 percent iron, is mined in open cuts by power shovels; but else- 
where underground mining is generally employed. 

These Mesabi range mines are about 1000 miles from coke- and 
steel-producing centers. In the metallurgy of iron "the ore generally 
comes to meet the coal." Why? At the mines, therefore, the ore is 
loaded into railroad cars of 60-ton capacity, and hauled to the Great 
Lakes port. Here, ore boats open their hatches, and a load of 10,000 
tons of ore is emptied into each boat in half an hour. After the long 
haul down the Great Lakes to lower lake ports, huge grab buckets 
scoop out the ore and drop it on a stock pile near the furnaces, 
which transform these mountains of ore into rails and other products. 

A smaller steel industry is located around Birmingham, Alabama, 
where a geological revolution once laid down coal, limestone, and 
iron ore. This region supplies about ten percent of our steel. 

Next in importance to the American iron ore deposits are those of 
the Ruhr and of Sweden, which also help to supply the furnaces of 
France, Germany, Belgium, and Luxembourg. The deposits in north- 
eastern England and Spain are valuable also. Russia, too, has many 
rich deposits of iron ore. Magnitogorsk, a city in the Urals, was 
built in 1928 by the Soviet Union around a mountain of a mag- 
netic iron ore called magnetite, Fe 3 O 4 . Many nations do not have 
deposits of iron ore. They depend for iron on other nations. 

What chemical reactions take place in the metallurgy of iron? 
The metallurgy of iron consists essentially of the reduction of iron 
oxide by means of carbon monoxide and heated coke. This reduc- 
tion takes place in a blast furnace. The charge that enters the furnace 
consists of hematite, coke, and flux. The hematite supplies the iron; 
the coke (and CO) reduces the hematite; and the flux removes the 
impurities by uniting with them, forming a molten slag which is 
drawn off. 

of Canada 

Surface mining of iron ore 
at Seven Islands, Labrador. 
The trucks haul the ore to 
a nearby railroad line to 
begin the long trip to blast 
furnaces in the United 




cold blast 

(hot gas) 

hot gas (part to heat stoves) 

Adapted from drawing by American Steel and Wire Company 

Fig. 90. Blast furnace and stoves. Study this diagram in conjunction with Fig. 91. 

The nature of the impurities in the ore determines the kind of flux 
used. Oxygen from the air that is forced through the charge keeps the 
coke near the bottom of the furnace actively burning. Equations for 
the reactions are given in Fig. 91. 

A blast furnace in operation. The modern blast furnace, which has 
been developed from a furnace invented five centuries ago, is cylin- 
drical, about 30 feet in diameter and 110 feet high, and is made of 
steel plates lined inside with firebrick, a refractory substance which 
withstands very high temperatures. At the top of each blast furnace 
is an apparatus through which the charge may be dropped with a 
minimum loss of heat. About eight feet from the base are 12 pipes, 
or tuyeres (twfrs) , through which 100 tons of hot air at about 
600C. are forced every hour. The heating of this air is carried out 
in the stoves, cylindrical towers 100 feet high. There are three or 
four stoves to each blast furnace. 

As reduction of the ore proceeds, the lava-like slag floats on the 
molten iron. The slag is removed about every four hours, just before 
the iron itself is tapped and poured from a taphole at the bottom of 
the furnace. The molten iron flows out along a gutter lined with 
firebrick until it reaches a brick-lined ladle capable of holding 100 
tons of metal. This is carried to an endless chain conveyer and 
poured into metal forms. The early forms into which the iron flowed 
resembled a litter of feeding pigs, henc% the/name pig iron. 

The blast furnace workj continuously except when repairs are 
necessary or business is slack. Often a blast furnace is in operation 
day and night for many months at a stretch. One blast furnace pro- 
duces between 600 and 1200 tons of pig iron daily. From the top of 



the furnace more than 3000 tons of hot gases issue every day. These 
gases, which may contain as much as 25 percent carbon monoxide, 
are gathered and used to heat the blasts of air that enter the tuyeres. 
Blast furnace gas is essentially the same as producer gas. 

Composition, properties, and uses of pig iron. Pig iron contains 
about four percent carbon and smaller amounts of compounds of 
sulfur, silicon, and phosphorus. These impurities make pig iron a 
grayish, brittle metal which melts at about 1150C. It cannot be 
forged or tempered, but can be cast into objects, such as sash weights, 
water pipes, stove parts, and radiators, that do not have to withstand 
great stresses or strains. 

Some pig iron is partially purified and changed to cast iron by 
melting it with scrap or wrought iron and then slowly cooling it. Its 

Fig. 91. Cross section of a blast furnace. Note the structure of the furnace, the 
nature of the charge and the chief chemical reactions which occur. 

Adapted from drawing by American Iron and Steel Institute 

waste gases 

to be used 

to heat air p^ I***|.J P^l ^ ore f char 9 e to be dropped 

for tuyeres 


lined with 

hot air 



taphole for pigiron 



carbon is present as globular particles of graphite. Such cast iron 
is less expensive than steel, is strong, machinable, and can be welded. 
It is used for tools, agricultural instruments, and machine parts that 
are not subjected to severe shocks. 

Most pig iron, however, is converted into steel by removing most 
of its impurities. In fact, in most plants, blast furnaces and steel 
converters stand side by side. 

Wrought iron oldest form of commercial iron. Up to the four- 
teenth century the form of iron most commonly used was wrought 
iron. This is almost pure iron, containing less than 0.1 percent car- 
bon. Wrought iron is soft and tough, and resists shocks and strains 
very well. Although it has been almost entirely replaced by soft 
steel, some of it is still used by blacksmiths, and very much more of 
it is used in the manufacture of chains, anchors, pipes, bolts, and 
temporary magnets. It resists atmospheric corrosion rather well. 

Wrought iron is prepared from pig iron in a shallow furnace hold- 
ing about 600 pounds of the melt. Iron oxide or rusty scrap iron sup- 
plies the oxygen that changes the carbon of the pig iron to carbon 

Fe 2 O 3 4- 3C -> 2Fe -f SCO f 

As the iron becomes purer, its melting point rises and the melt be- 
comes a pasty mass. This is stirred, puddled, and then removed in the 
form of fiery balls at the end of iron rods. These balls are squeezed 
between hammers, which remove most of the slag, and then rolled 
to distribute the remaining slag throughout the iron, giving the 
wrought iron a fibrous structure. 

Henry Bessemer and William Kelly usher in the steel age. Both 
wrought iron and steel have been known for centuries. They were 
made in small amounts by skilled workers using a slow and expen- 
sive process until the invention of the rifled cannon, locomotive, and 

Adapted from drawing by American Iron and Steel Institute 

*lfev ;! /" co ' 


lining of 
sand bri 



Fig. 92. Cross section of a 
Bessemer converter. Note the 
chemical changes which oc- 
A cur as the air passes through 
slag ^e molten iron. 

pig iron 


other machines brought an unprecedented and immediate demand 
for huge quantities of steel. 

William Kelly, a Pittsburgher living in Kentucky, got the notion 
that a blast of air would not chill molten pig iron but would increase 
its heat by oxidizing its impurities. In this way he planned to make 
steel quickly and inexpensively. He was scoffed at, and his father-in- 
law actually questioned his sanity and had a doctor examine him. 

Kelly went right ahead, however, and in 1851 he gave a public 
demonstration of his process. An eye-witness reported: "We saw 
a vessel that had a mouth open on one side and near the top. The 
whole was shaped something like an egg. We saw molten metal 
poured into the vessel. Then Kelly turned on a blast of cold air. The 
vessel set up a roaring noise, and fire belched furiously from its 
mouth, making many colors. But only for a few minutes. The noise 
and fire died down. We then saw a blacksmith take a small part of 
the iron which had cooled, and contrived and threw at the feet of 
the amazed spectators a perfect horseshoe. No one laughed at Kelly 

While Kelly worked to improve his product, Henry Bessemer in 
England was discovering the same process independently. During 
the Crimean War, he had invented a new type of cannon but he 
could not find an iron strong enough to withstand the high pressures 
of the expanding gases released. This led him to researches in steel, 
and in 1855, at the age of 42, he solved the problem and obtained 
a patent on his process. Kelly then quickly obtained a patent for his 
own process in the United States. Bessemer built a more efficient 
furnace and bought out Kelly's patent. 

Bessemer was later knighted. This honor came to him not for his 
invention of the steel furnace but rather for a suggestion he had 
made when he was 20 years old. He had recommended an improve- 
ment for preventing the re-use and counterfeiting of seals and stamps 
on official documents. 

At this time steel was selling at $250 a ton. The world was wait- 
ing for an inexpensive steel, and, almost overnight, the iron indus- 
try was revolutionized. The first Bessemer furnace, or converter, 
used in the United States was set up at Troy, New York, in 1864. 
The new process did in ten minutes what the old process took a 
whole month to accomplish. The old puddling furnaces that pro- 
duced wrought iron were almost completely scrapped and soon 
steel took the place of wrought iron almost entirely. 

In production, pounds were replaced by tons. By 1870 steel pro- 
duction had forged ahead so rapidly that it equalled the production 

One of the most spectacular dis- 
plays in industry a Bessemer 
converter during a "blow." 

Kcthkhcm Steel Company 

of wrought iron. At present, wrought iron represents only about one 
percent of the total output of pure iron in this country. The highly 
skilled ironworker with his small furnace has given way to hundreds 
of trained men handling huge furnaces turning out steel on a gigan- 
tic mass-production scale. 

The acid Bessemer process, a fiery spectacle. The present-day Bes- 
semer converter used in the United States is a barrel-shaped pot of 
wrought steel about ten feet in diameter and 20 feet deep, lined to 
a thickness of 1.5 feet with heat-resisting brick made up largely of 
sand, SiO,, an acid anhydride. Fifteen tons of a fiery broth of mol- 
ten pig iron, whose chief impurities are carbon and silicon, are 
poured into the Bessemer converter from the blast furnace. 

Through this molten mass blasts of cold air are forced under about 
25 pounds pressure, through holes at the bottom of the converter. 
This blow lasts between ten and 20 minutes. The oxygen in the blasts 
of air oxidizes the iron it strikes, first forming ferrous oxide, which 
in turn reacts with the silicon impurity, forming silicon dioxide and 

2Fe + O 2 - 2FeO 
Si + 2FeO -> SiO 2 + 2Fe 

Silicon dioxide, insoluble in the molten iron, accumulates as slag. 

As the blow continues, the carbon impurities begin to burn. A 
roaring boil then takes place in the converter. The carbon monoxide 



burns at the mouth of the converter, and countless flying sparks of 
metal and slag add to the spectacle. 

C + O > CO 2 
C0 2 + C -) 2CO 
2CO + 2 - 2CO 2 

In a few minutes the carbon is gone; the (lame flickers and contracts. 
This is the signal to stop the blast by turning the converter vessel 
over on its side. 

Inside the converter, a seething mass of molten iron is covered with 
a thin layer of slag. The metal, however, contains some dissolved 
gases and, if solidified at this point, would be spongy. So a prede- 
termined amount of carbon and manganese is added in the form of 
iron alloys, such as ferromanganese, which contains about 70 percent 
manganese, or spiegeleisen^ which contains about 20 percent man- 
ganese. About 14 pounds of manganese are used for every ton of 
steel produced. The manganese unites with any dissolved air or com- 
bined oxygen, and in addition strengthens the steel by its presence. 

The chemical reactions that take place in the converter when 
spiegeleisen is added are: 

<\ r .. r^u f Mn + 2FeO > MnO 2 + 2Fe 

1) The action of the manganese: | Mn + ^ _ MnQ * 

2) The action of the added carbon: 

3Fe -f C > Fe 3 C (iron carbide) 

The converter is then tipped completely over and its charge of liquid 
steel is poured into a waiting mold. 

The basic open-hearth process. Pig iron containing sulfur and 
phosphorus cannot be treated in the acid Bessemer converter, since 
this type of furnace cannot remove these impurities. Phosphorus is 
oxidized during the blow, but it will not enter the slag and become 
part of it. 

The basic open-hearth furnace, however, completely removes both 
sulfur and phosphorus (acid impurities) by means of lime and a 
lining made of the oxides of calcium and magnesium which are basic 
anhydrides. In England and on the Continent the Bessemer furnace 
lined with these basic oxides is used in the treatment of iron ores 
containing sulfur and phosphorus. 

The open-hearth furnace was invented by Charles W. Siemens 
(ze'mens) , who left Germany to become a citizen of England. He 
patented his furnace five years after the Bessemer furnace was in- 
vented. In 1873 the first open-hearth furnace was built in the United 



Fig. 93. Open-hearth furnace 
(greatly simplified) showing 
the nature of the charge and 
lining and the flow of the hot 

The present open-hearth furnace, as large as an eight-room house, 
contains a shallow, obiong steel basin about 50 feet long by 15 feet 
wide encased in brickwork. The charge consists of about equal 
amounts of scrap iron and molten pig iron, with a small amount of 
iron ore placed over a layer of limestone, CaCO : ,. The scrap iron con- 
sists of old, discarded factory and farm machinery, junked cars, and 
so forth. This is cheaper than pig iron, contains less carbon, and 
conserves our reserves of raw material. Burning fuel and hot air 
enter at one end, are deflected down from the low roof made of re- 
fractory brick, and heat the charge just as a gigantic blast lamp 
might; the waste gases finally leave at the other end of the furnace. 
Such a furnace in which the charge is heated by flames deflected 
from the roof is called a reverberatory furnace. By a regenerative 
process, all the heat of the gases is used by reversing their flow at 
regular intervals. The following reactions take place: 

1) Carbon is burned out by the oxygen of the air 

and by the oxygen of the rusty scrap iron, and ore. 
3C + Fe 2 O 3 ~> 3CO + 2Fe 

2) The limestone decomposes, forming carbon dioxide, which 
bubbles up rapidly through the melt. This bubbling produces a 
vigorous stirring and leaves a basic lining of CaO. 

CaC0 3 - CaO + CO 2 1 

The open-hearth depart- 
ment of a steel mill. Open- 
hearth furnaces, located 
behind the wall at the right, 
are charged with molten 
iron from the blast furnace, 
scrap iron, and limestone. 
About 12 hours are re- 
quired to complete a "heat" 
of 250 tons of steel. 

A special machine for the man- 
ufacture of seamless steel tubing. 
Shown is a length of white-hot 
tubing just after piercing. During 
the manufacturing process a steel 
billet eight feet long is converted 
to a tube 35 feef long. 

Rcpuhlic Stcd Corporation 

3) Some sulfur is liberated as sulfur dioxide. 
' '-' < v - : ; - " FeS + 2FeO -* 3Fe + SO 2 1 

4) The calcium oxide combines with the phosphorus and silicon 
oxides, forming a slag, which rises to the top. 

3CaO 4- P 2 O 5 -> Ca,(PO 4 ) 2 
CaO H- SiO 2 -> CaSiO 3 

From time to time, samples of the 200- to 500-ton pool of liquid 
are removed and quickly analyzed for carbon content so that the 
composition of the steel may be controlled carefully. This open- 
hearth process is longer and may be controlled more easily than the 
acid Bessemer process, and takes about 12 hours for 200-ton heats. 

In spite of its slightly higher price, basic open-hearth steel is in 
great demand in this country because of its generally higher uni- 
formity and quality. In fact, less than five percent of the steel made 
in the United States today is prepared by the acid Bessemer process. 
Open-hearth steel is used for rails, structural steel, machinery, and 
plates for boilers and ships; Bessemer steel for sheets, pipes, and bars. 

From converter to rail, girder, sheet, and pipe. When the steel 
is ready to be tapped, a hole is opened in the rear of the furnace, and 
the steel drains into a huge brick-lined ladle. The lighter slag flows 
out last and overruns, or overflows, the filled ladle. Between five 
and 15 tons of steel from this ladle are poured into square, round, 
or rectangular cast iron molds called ingot molds. The molds are 



stripped from the ingots, which are then lowered into underground 
furnaces or soaking pits, where they are heated for six hours to a 
uniform rolling temperature of about 2400 F. 

All steel begins in the ingot. While still red hot, the ingot, as tall 
as a man, is first descanted, that is, cleaned of all surface scale and 
seams. Then it is changed into smaller sizes and different shapes in 
rolling mills or under hammers and presses. In the continuous-sheet 
rolling mill slabs of hot or cold steel are rolled down to specified 
widths and thicknesses in one uninterrupted operation at a rate of 
about 1000 feet per minute. In the blooming mills, a pair of heavy 
steel rolls, built like a huge clothes wringer, flatten and shape the 
ingot into a bloom, which may then be converted into a rail. The rail 
travels on another roller line to a hot saiv, where it is cut into 33-foot- 
6^-inch lengths. When cold, each piece becomes a standard rail 33 
feet long. Steel rails at the mill cost about 4.5 cents a pound. Sim- 
ilarly, other steel ingots are sent from one machine to another and 
finally emerge as rods, bars, plates, girders, wire, and tubing. 

Composition and properties of steel. Steel contains from 0.05 per- 
cent to 1.7 percent carbon, the percentage being regulated according 
to the use to which the steel is to be put. Most steel does not con- 
tain sulfur or phosphorus, for sulfur makes the steel brittle when hot, 
and phosphorus makes it brittle when cold. Steel is hard, silvery in 
color, ductile, magnetic, and can be welded, forged, and tempered. 

Tempering is one way to change the properties of steel. The per- 
centage of carbon in steel and the form in which the carbon is pres- 
ent determine the properties of the steel. The carbon may be present 
as free carbon or as iron carbide, Fe 3 C. It may be present either in 
solution or as crystals. A steel rich in dissolved iron carbide is very 
hard and brittle, while one containing a large amount of crystalline 
iron carbide is comparatively soft. 

The percentage of free and dissolved iron carbide in steel may be 
varied by heat treatment. This treatment of a metal to vary its hard- 
ness, strength, and toughness is called tempering. Rapid cooling or 
quenching in cold water or brine produces a hard but brittle steel. 
By quenching in boiling water or hot oil, a slower rate of cqpjing is 
obtained, resulting in partial crystallization in the hard sffe^l and 
producing a soft but tough steel. 

Heat without flame, and almost without wires. Another method 
for heat-treating steel is by means of an induction furnace. No flame 
of any kind is involved, nor are there direct connections to a source 
of electricity. Moreover, instead of being heated chiefly on the sur- 
face, as in most kinds of heating, a substance placed in an induction 



furnace is heated from within. Thus, the heating is very rapid, easily 
controlled, and applied only where it is needed. A small induction 
furnace using 600 kilowatt hours of electricity is capable of melting 
one ton of steel per hour. 

An induction furnace acts as a transformer. Alternating current 
of high frequency flows in the primary coil, and the material to be 
heated is placed in much the same position as if it were the core of a 
transformer within the secondary coil. As the transformer operates, 
the core becomes warm. If the frequency of the current fed into the 
primary coil is increased greatly, the core becomes hot. This heating 
effect comes from eddy currents set up within the core as a result 
of the changing magnetic fields produced by the rapidly alternating 

Steel alloys, another way to change the properties of steel. Heat- 
treating is not the only method of improving steel. The addition 
of small amounts of other metals changes the properties of steel and 
widens the range of uses of the alloy steels. It must be remembered 
that steel itself is an alloy of iron and carbon. 

With the development of the railroad, automobile, and the air- 
plane, engineers raised a cry for more alloys. Transportation needed 
strong, flexible, rust-resisting alloys. Out of the chemist's crucible 
came hundreds of such new alloys for thousands of different uses. 
On the wings of this new age of alloys came molybdenum and vana- 
dium steels for axles, springs, gears, and rails; wolfram, molybdenum, 
arid chromium steels for high-speed cutting tools; nickel-chrome 
steels for armor plate; cobalt steel for turbine blades, tail cones, and 
combustion chamber linings of jet planes; and manganese steel for 
brake shoes, burglar-proof safes, the jaws of rock crushers, and non- 
magnetic binnacles of ships. 

A far-reaching development in steels is the large-scale production 
of stainless steels. These alloys, first used for making stainless kitchen 
utensils, are employed in a multitude of ways. For example, the Em- 
pire State and Chrysler buildings in New York City are sheathed in 

Adapted from a drawing by American Iron and Steel Institute 

Fig. 94. Cross section of an 
electric furnace for the pro- 
duction of alloy steels. The 
chemical reactions involved 
are essentially the same as 
those of the open-hearth 


top removes 
for charging 
scrap -x^ 







rustless steel that needs neither paint nor any other coating to pro- 
tect it against corrosion. 

A typical stainless steel is "USS 18-8" which, besides iron, con- 
tains about 18 percent chromium and eight percent nickel as well 
as 0.5 percent manganese, 0.5 percent silicon, and about 0.1 percent 
carbon. The trademarks of two other stainless steels are "Enduro 
KA2" and "Rezistal." The metal niobium is also added to stainless 
steel. Small amounts of silver added to stainless steel decrease salt- 
water corrosion of this alloy. "Duriron," which resists acid corrosion 
very well, contains about 15 percent silicon. "USS Cor-Ten" contains 
copper and is a very corrosion-resistant steel used widely in railroad 
passenger cars. 

Some high-grade steel, used for special purposes such as razors, 
springs, and pens, is made in small amounts in electric arc furnaces 
from ingot iron and steel plus the necessary alloying elements. The 
electricity serves to supply the high temperature required for the 
purification of this fine steel. The process can be rigidly regulated. 

What is case-hardening? In the process called carburizing, wrought 
iron or soft steel is heated in contact with coke or charcoal. The car- 
bon forms a layer of very hard, high-carbon steel over the softer low- 
carbon metal. This gives a hard skin or exterior and a tough interior 
a condition ideal for gears and carving knives. This method was 
the one used by the skilled artisans who fashioned swords, armor, 
and guns from wrought iron hundreds of years before Bessemer steel 
came into use. 

Case-hardening is also accomplished by the use of ammonia and 
sodium cyanide. In this process, called nitriding, ammonia gas at high 
temperatures or NaCN reacts with special steels in airtight contain- 
ers. Very hard nitrides are formed on the surface of the steel. 

Bell Telephone Labomto 

Tiny flaws and defects in 
metal are detected by this 
instrument called a Met al- 
lograph. The Metallograph 
takes photomicrographs of 
the surface of a metal mag- 
nified up to 2,000 times. 

A'-/,',,.!/ i 

A radiograph of a weld shows a crack (left) in the internal weld 
metal. The numerals are used to identify the location of defects. 

How steel is analyzed. Even traces of impurities or slight defects 
in the inner structure of steel alloys may weaken them. Many rail- 
road wrecks and explosions have been laid at the door of defective 
steel. Great strides have been made in the analysis of the structure 
of steel. Metallography is the name given to this field of study. Vari- 
ous methods of testing steels are used. Besides testing the physical 
properties of steel, including strength, ductility, and hardness, accu- 
rate chemical analyses give its composition. 

By etching the carefully polished surface of a sample of an alloy 
with a dilute acid, the irregularly pitted surface, photographed un- 
der a special microscope, will show poor welds, cracks, blowholes, 
sandholes, and any other weaknesses. Whether the alloying metal has 
formed a compound or is merely in solution can also be determined 
from these photomicrographs. 

X-ray photography is also used in the examination of steel. Port- 
able x-ray machines have been developed that are powerful enough 
to affect a photographic film behind solid steel one foot thick. Struc- 
tural- defects in castings or welds arc thus easily detected. Special rail- 
way cars carrying x-ray machines and other equipment examine thou- 
sands of miles of railroad track each year. 

New horizons in the steel industry. Several major innovations are 
now taking shape in the steel industry. 

First, in an effort to raise steel production with present plants and 
equipment, air enriched with pure oxygen is being used instead of 
ordinary air in a few blast furnaces and steel converters. Oxygen- 
enriched air in the blast furnace hastens the melt by raising the tem- 
perature several hundred degrees. In the open-hearth and Bessemer 
furnaces, the addition of pure oxygen to the air shortens the produc- 
tion time by accelerating the removal of carbon. Cheaper oxygen may 
accelerate this new development. 

Second, high pressure operation in the blast furnace saves time and 
increases the daily output of pig iron. By means of a carefully worked- 
out system of valves, the pressure inside the blast furnace can be 
carefully increased and regulated so that the rate of the blast can be 
stepped up. This results in increasing the daily output of pig iron 
and also reduces the consumption of coke by about ten percent. 



Use by approximate 

Building and 

Industrial machinery 
and equipment 

Petroleum, gas 
and mining 

Steel converting 

Domestic and 
commercial appliances 
and utensils 

Agricultural machinery 
and equipment 

Export and other 

Third, success has already crowned the efforts of one very large 
steel company in this country in its long attempt to substitute a 
quicker, cheaper, and more compact method of converting the molten 
steel of the converter directly into semi-finished steel products. The 
method, announced in 1948, is known as the continuous casting proc- 
ess. This process eliminates the most expensive and intricate ma- 
chinery now used in processing steel. All finished steel products 
formerly began as ingots. However, continuous casting makes use of 
a relatively small casting machine which forms billets of iron, cools 
them, and then cuts them to any desired length up to 35 feet. It 
eliminates hammering, pressing, rolling, and reheating. Not only does 
this process cut costs and increase production, it also makes possible 
smaller and less costly plants. In time this development may help de- 
centralize the steel industry and allow even greater expansion. 

Finally, work is progressing to meet the ever-increasing threat of 
a shortage of top-grade iron ore containing 60-65 percent iron. With 
the already rapid depletion of this rich iron ore of the Mesabi range 
mines, iron men are studying how to concentrate and utilize low- 
grade iron ores containing only 25 percent iron found plentifully 
in many parts of the United States. For example, northern New York 
iron mines, abandoned almost half a century ago, are being reexam- 
ined. Inferior iron ores (magnetic taconite) from the Mesabi range, 
have been ground, impurities removed magnetically or by flotation, 
heated into walnut-sized pellets containing 65 percent iron, and then 
used in blast furnaces. If successful, such benefication of low-grade 
ores will give us a plentiful supply of iron foj nriny years to come. 
In the meantime, we have begun to import rich iron ores contain- 
ing as high as 67 percent iron from Labrador, Venezuel^, Brazil, The 
Dominican Republic, and even Liberia. s 


de Kruif, Paul. Seven Iron Men. Harcourt, Brace 8c Co., 
New York, 1929. The author of Microbe Hunters tells the 
story of the iron-hunting Merritt family, pioneers in the Min- 
nesota iron industry. 

Lyans, R. C. "Student-Made Visual Aids." Journal of Chem- 
ical Education, May, 1944, pp. 241-242. Includes the blast and 
open-hearth furnaces. 

Morrow, Martha G. "Steel for War and Peace." Chemistry, 
April, 1945, pp. 1-8. Well written and beautifully illustrated. 

Parker, Charles M. "Panorama of Steel." Journal of Chem- 
ical Education, May, 1951, pp. 236-242. An illustrated article 
on recent developments in steel. 


Group A 

1. (a) Where does Fe rank with respect to usefulness 
among the metals? (b) Where does Fe rank with respect to 
tonnage used? 

2. (a) Where and (b) in what form does Fe occur in the 
United States? (c) in Europe? 

3. (a) Why are the mines of the Mesabi range so valuable? 
(b) How is this ore mined? (c) Where are the reduction fur- 
naces? (d) Why? 

4. Make a diagram of a blast furnace and label: the place 
where the charge enters, the tuyeres, the hotter parts of the 
furnace, the cooler parts, the place where the slag collects and 
is tapped, and the place where the Fe collects and is tapped. 




5. (a) Of what is the solid charge in the blast furnace com- 
posed? (b) What is the function of each constituent of the 
charge? (c) What other substance must be used in the furnace 
in large quantities? (d) For what purpose? 

6. (a) Describe what takes place in a blast furnace, 
(b) Write equations for the chemical reactions that occur. 

7. (a) What valuable gaseous product is formed during 
the operation of the blast furnace? (b) For what is it used? 

*== ._i_ *===* 

8. Copy and complete. Do not write in this book. Blast fur- 
nace iron is also called . . . iron. It contains about . . . percent 
of ... and smaller amounts of compounds of . . . , . . . , and .... 
The impurities of the pig iron are removed by either the . . . 
process or the . . . process. 

9. (a) What properties of pig iron make it useless for many 
of the purposes for which steel is made? (b) For what is pig 
iron used? 

10. (a) Name two men who helped usher in the steel age. 
(b) At what important period of American history did the steel 
age begin? 

11. Using a labeled diagram, describe the acid Bessemer 

12. (a) Why must a Bessemer converter tilt easily? (b) How 
does the operator know when to tilt it? 


13. What is the source of the highest temperature reached 
by the contents of the converter? 

14. (a) How are the impurities of pig iron in the Bessemer 
converter removed? (b) Write equations for the reactions. 

15. A function of the Bessemer process is to remove almost 
all C, yet C is added at the end of the process. Explain. 

16. What is the function of spiegeleisen in the operation of 
the Bessemer converter? 

17. The acid Bessemer process removes all objectionable 
impurities from only one kind of ore. Explain. 

18. Make a drawing of the basic open-hearth furnace. Label: 
the layer of slag, the place where the molten steel collects, the 
place from which the steel drops into the ladle. Show by arrows 
the path of the air and the combustible gases. 


19. (a) Why is the open-hearth process called basic? 
(b) What is the source of the great heat produced? 

20. (a) Why is the basic open-hearth process said to be re- 
generative? (b) Why is the furnace called a reverberatory 

21. (a) Of what does the charge in the basic open-hearth 
process consist? (b) Why is scrap iron used? 

22. Show, by equations, the three most important chemical 
changes that take place in the basic open-hearth furnace. 

23. Compare acid Bessemer and basic open-hearth steel by 
making a table listing (a) quality, (b) cost, (c) uses, (d) ton- 
nage, (e) time required for the process. 

24. Compare the properties and composition of steel and 
pig iron. 

25. Why must most steel be free from S and P? 

26. Name two general methods for changing the properties 
of steel. 

27. (a) What is tempering? (b) How is it done? 

28. What physical and chemical changes take place in steel 
during tempering? 

29. (a) What is an alloy? (b) How does it differ from a 

30. Steel itself is an alloy, (a) Explain, (b) Name four steel 
alloys and (c) tell for what use each is suited. 

31. How is steel made in electric furnaces? 

32. A sample of hematite contains 75 percent Fe 2 O 3 . What 
weight of wrought iron can be made from 250,000 tons of this 

33. What volume of air is required to oxidize all the C 
present in 500 tons of pig iron containing four percent C? 

34. What are three tests used in testing the quality of a 
sample of steel? 

35. (a) What is the oldest form of commercial Fe? (b) How 
does it compare in composition and properties with steel? 

36. What is the chief chemical reaction that takes place 
during the making of wrought iron? 


37. Why does the wrought iron in the puddling furnace be- 
come pasty? 

38. How do you account for the fact that fairly pure Fe has 
been known for 5000 years, although Fe is never found free? 

39. Why does the manufacture of basic open-hearth steel 
outrun Bessemer steel about 13 to one? 

40. When would the addition of flux in a blast furnace be 

41. What is the difference between an acid lining and a 
basic lining? 

42. (a) State four new developments in the steel indus- 
try, (b) Describe one of them in detail. 

Group B 

43. Of what use is slag from a blast furnace? 

44. (a) Name and give the formulas of three ores of iron 
other than hematite, (b) What is fool's gold? 

45. 19 g. of Fe yielded 27.2 g. of Fe 2 O 3 . Calculate the atomic 
weight of Fe. 

46. What four factors help to determine the location of 
plants for the production of iron and steel? 

47. (a) Give the names and formulas of two iron oxides 
used as pigments in paints, (b) What property makes them 
ideal, for this purpose? 

48. Why does the molten metal not run out of the holes in 
the false bottom of the Bessemer converter? 

49. Describe the induction furnace and its operation and 

50. What is the difference between carburizing and ni- 


1. Make a list of the different metals that enter into the 
manufacture of an automobile. From this list pick out those 
metals that must be imported. Why must these metals be im- 
ported? Of what significance is this fact? 

2. On an outline map of the'United States indicate the lo- 
cation of the largest steel mills and the richest coal deposits 
in this country. Show also the location of most of the iron 
mines and the routes over which both coal and iron are sent 
to the steel plants. What relationships do you discover? 

3. Construct a model of an open hearth or blast furnace. 



. . . Berthelot analyzed a small vo- 
tive figure from the excavations of 
ancient Chaldea and found it to con- 
sist of nearly pure copper. The age 
of this figure is variously estimated 
at from 3000 to 4000 B.C. A small 
metal cylinder from Egypt of a period 
estimated at about 4000 B.C. was 
also of copper. Thus the mining and 
metallurgy of copper is at least 5000 
years old. J. M. Stillman, 1919 

Copper, the most important nonferrous metal. In 1831 Michael 
Faraday performed an immortal experiment. By moving a magnet 
through a coil of copper wire, he discovered that an electric current 
could be produced in a wire. This was the beginning of the electric 
age, with its generators, motors, electric wires, cables, and thousands 
of electric instruments. The manufacture of all copper items in the 
United States consumes over one million tons of copper each year. 

If steel is the backbone of the machine age, then copper is its nerv- 
ous system, for along copper wires travel the electron impulses that 
animate our huge machines. Next to silver, copper is the best con- 
ductor of electricity. 

Copper, the red metal, was the first metal used by man. It was 
used during the Bronze Age long before iron, mainly because cop- 
per was found free. Copper is soft, has a high tensile strength, and 
melts at about 1100C. The ancients used copper and its alloys for 
implements and decorations. 

Copper "nerves" carry electricity throughout the world. Copper is 
such an excellent conductor of electricity that millions of miles of 
copper wire are used to carry electric power. Great quantities of 
copper are used also in the construction of telephone and telegraph 



Use by approximate 

Industrial machinery 
and appliances 

Heating and 
plumbing equipment 

Communications and 
transmission, etc. 

Hardware, tools, utensils 

wires and cables and in underwater cables for use in transoceanic 
communication. Parts for electric motors, electronic devices, and 
other equipment using electricity require great quantities of copper. 

Alloys of copper. Bronze, composed of copper, zinc, arid tin, and 
brass, composed of copper and zinc, are among the most widely used 
alloys of copper. Nickel silver, or German silver, is an alloy of cop- 
per, zinc, and nickel. "Monel metal" is composed of copper, nickel, 
and iron. This alloy is widely used in the production of noncor- 
rosive sinks, cooking utensils, and food and chemical processing 

Copper is used in varying amounts in the production of many of 
the alloys of steel, aluminum, and magnesium. Our one-cent piece is 
made of a copper alloy, and other coins contain copper. 

Copper in construction. Because copper and its alloys form self- 
protecting films on exposure to moist air (see page 431) , and are of 
relatively high tensile strength, they are widely used in the construc- 
tion industries. Gutters, downspouts, valleys, and flashings made of 
copper or copper alloys usually last as long as the house or other 
building on which they are installed and require a minimum of 
maintenance. Copper sheathings and copper roofs are also used. Cop- 
per sheathing frequently is used as a covering to protect the hulls of 
wooden ships. ' * 

Brass pipe and copper tubing are often specified by architects in 
plumbing and heating installations; and faucets, valvefc, and other 
plumbing fixtures are usually made of brass. Hardware and screens 


Electrical machinery 
and apparatus 

Metal stampings and 
wire products 

made of copper alloys resist corrosion and last longer than those 
made of most lower cost metals. 

Copper and the processing industries. The processing industries 
convert raw materials into finished or semifinished products. Many of 
the processes involved are chemical in nature and must be carried out 
in vats, reaction chambers, cookers, and similar equipment that will 
withstand, without corrosion, the chemical reactions, the high tem- 
peratures and pressures, and various other factors that may be pres- 
ent. Copper and copper alloys are often ideally suited to the require- 
ments of a particular processing situation, and great quantities are 
used in the production of processing equipment. Copper and copper 
alloys are used also in the production of household utensils. Copper- 
bottomed stainless-steel cooking utensils are very popular. 

How copper is freed from its ores. When copper occurs free, as it 
does in the Lake Superior region, its ore is melted, and the copper 
flows away from its impurities. In the Lake Superior region, most of 
the copper ore occurs as small particles of free copper embedded 
in gangue, from which it must be separated. 

Though lumps of native copper weighing 500 pounds have been 
found, the metal occurs chiefly in the combined state. Formerly, 
mines in the Keweenaw Peninsula of Michigan supplied much of our 
domestic copper, but now most of it is mined in Arizona, Montana, 
Nevada, New Mexico, and Utah. In these western areas copper usu- 
ally occurs as the sulfide, Cu 2 S, together with iron sulfide and copper 


In an underground Montana 
copper mine, a miner drills 
blast holes with a pneumatic 


ts Research Association 

The metallurgy of sulfide ores is very complex. It involves lirst the 
process of roasting, during which the crushed ore is heated in the 
presence of air and changed into cuprous oxide. 

2Cu 2 S + 30 2 -> 2Cu 2 + 2S0 2 1 " ** 

This cuprous oxide is then transferred to a reverberatory furnace 
where coke or powdered charcoal reduces the copper oxide to copper. 

Cu 2 O + C -> 2Cu + CO T 

Matte from blast or reverberatory furnaces contains copper and 
iron sul fides. The matte is treated in a furnace resembling the Besse- 
mer converter, where oxygen from blasts of air unites with the sulfur 

An open-pit copper mine in New Mexico. 

r () />/i,r mnl Rrn*x Itwnrrh Awarifitinn " '"i" 



in the matte, forming sulfur dioxide. The iron oxide remaining in 
the furnace combines with the silicon dioxide lining of the furnace, 
forming ferrous silicate slag, FeSiO. $ . The copper that results from 
this process is called blister copper, because it has cavities and other 
defects caused by escaping gases. Blister copper may contain as much 
as 99 percent copper. 

How blister copper is refined electrolytically. Most copper ores 
contain small amounts of silver, gold, zinc, arsenic, and other metals. 
Since the presence of minute amounts of certain of these substances 
greatly decreases the ability of the copper to conduct electricity (0.03 
percent of arsenic lowers it 14 percent) , blister copper is purified 
electrolytically. Plates of blister copper about three feet square and 
one inch thick are suspended as anodes in tanks containing a warm 
solution of copper sulfate and sulfuric acid, which is stirred mechan- 
ically. The cathodes consist of very thin sheets of pure copper. 

Impure copper passes into solution at the anode, producing cop- 
per ions. The two electrons thus released reach the cathode, where 
they join a copper ion and convert it into free copper, which is 

Pouring molten copper into wire 
bars from which electric wires 
are drawn. 



deposited on the cathode. Negative sulfate ions, under the influ- 
ence of a direct current, travel to the anode, give up their extra 
electrons, and unite with the copper of the anode, forming more 
copper sulfate. Thus, while copper from the anode is gradually 
transferred to the cathode, the concentration of the copper sulfate 
bath does not change (see Fig. 95) . 

Cu - 2e -> Cu++; Cu++ + 2e -> Cu 

Because copper is higher than either silver or gold in the replace- 
ment series, it loses electrons more easily than either of these rarer 
metals. Hence, while copper goes into solution and is carried to the 
cathode, the silver and gold do not ionize but sink to the bottom of 
the tank as sludge. Zinc goes into solution, and arsenic also collects 
in this sludge. These metals are present in much smaller amounts 
than copper and require a higher voltage to reduce their ions and 

A copper converter in which matte is reduced 
to form slightly impure, or blister, copper. 

Fig. 95. Electrolytic process for 
refining copper. How should 
this diagram be altered to 
illustrate the copper-plating 


Sludge of 

precipitate them out on the copper cathode. Huge amounts of gold 
and silver are later extracted from the sludge. 

The principle involved in this electrorefining of copper is used 
also in copper-plating. The object to be plated is hung from the 
cathode; the anode consists of a slab of pure copper. As dissociation 
occurs, a coating of copper gradually builds up on the cathode. 

Flotation, a widely used ore-dressing process. One of the most im- 
portant metallurgical operations introduced in recent years in ore 
dressing is flotation, which was developed in an attempt to concen- 
trate copper ore and get rid of most of the gangue before the ore is 
chemically treated to obtain its copper. Ores containing as little as 
one percent copper are processed in the United States. Flotation 
makes possible economical mining of such ores. 

In the flotation process, the ore is crushed and placed in a tank 
containing water and two chemical agents. The first of these is a 
frothing agent, such as pine oil, the purpose of which is to create 
oil -coated air bubbles, when air is blown into the tank. The second 
chemical is a collecting agent which forms a water repellent film 
around the particles of metal. The particles attach themselves to the 

Adapted from a drawing of Denver Equipment Company 

Fig. 96. Cross section of an 
re-flotation tank in which 
the crushed ore, air, water, 
and chemical agents are 


Copper and Brass I{<tirrli Association 

An ore-flotation tank in operation. 

air bubbles and ride to the surface. The froth is then skimmed oft 
and treated chemically to obtain its copper. 

One of the pioneers in the discovery of the "affinity of oils and 
fatty substances for mineral particles" was an American woman, Mrs. 
Carrie Billings Everson. She patented this process in 1886. 

Chemical properties of copper. Copper occurs just below hydrogen 
in the replacement series. Therefore, it loses electrons less easily 
than iron, zinc, tin, and the other metals above hydrogen, but more 
easily than silver, mercury, and gold. Therefore a piece of iron placed 
in a solution of copper sulfate becomes coated with copper. The 
copper takes the electrons lost by the atoms of iron which have gone 
into solution. A strip of silver, on the other hand, when placed in 
a solution of copper sulfate, remains unchanged. 

Cu++ + SO 4 + Fe - 
Cu++ 4- SO 4 - 4- Ag - 

> Fe++ + SO< ' + Cu 

> no reaction 

This is made use of in the recovery of thousands of tons of copper 
each year. Water from the mines containing CuSO 4 is passed over 
scrap iron which is replaced by pure copper. 



Dry air hardly affects copper. But when copper is exposed to 
moist air, it first becomes coated with a thin layer of red cuprous 
oxide, Cu 2 O. The water vapor and carbon dioxide of the air react 
with this Cu 2 O, forming a dull greenish layer of basic copper car- 
bonate, CuCO 3 Cu (OH) 2 . The greenish coating, or patina, on cop- 
per roofs, copper pipes, and bronze statues is generally considered to 
be a basic copper carbonate accounted for by this chemical change 
and a basic copper sulfate. Both copper oxide and basic copper car- 
bonate act as protective films that prevent deeper corrosion. 

Copper unites with nearly all the nonmetals, especially at high 
temperatures, forming copper salts. Thus, heated copper burns in 
an atmosphere of chlorine, forming thick clouds of copper chloride. 

Cu + C1 2 -> CuCl 2 

Since copper occurs below hydrogen in the replacement series, 
it cannot replace hydrogen in acids. For this reason, hydrochloric 
acid and dilute sulfuric acid do not attack copper. The oxidizing 
acids, nitric and concentrated sulfuric, react with copper, yielding 
not hydrogen, but reduction products of the acids, thus: 

3Cu + 8HNO 3 -> 3Cu(NO 3 ) 2 + 4H 2 O + 2NO t 
Cu + 2H 2 SO 4 (hot cone.) -> CuSO 4 + 2H 2 O + SO 2 1 

Why copper exhibits two valences. Copper has two valences, 
plus one and plus two. The electron theory offers the following in- 
teresting explanation of these two valences: in the atom of a heavier 
element, the electrons are not always arranged in the manner we 
would expect from its atomic number. For example, in the case of 
copper (atomic number 29) , the 29 electrons outside the nucleus 
may be distributed as follows: 

Cu: Valence of 1 Cu: Valence of 2 

2 electrons in first ring 2 electrons in first ring 

8 electrons in second ring 8 electrons in second ring 

18 electrons in third ring 17 electrons in third ring 
1 electron in last ring 2 electrons in last ring 

This arrangement occurs because the heavier elements, having 
larger nuclei, have the power to stabilize more than the normal 
number of electrons in their inner rings (see page 160) . Iron, for 
example, has valences of two and three, and mercury has valences of 
one and two. * 

Because copper has two different valences, it forms two series of 
salts. Cuprous salts contain monovalent copper; cupric salts contain 


divalent copper. The cuprous ion is almost colorless; the cupric ion 
is bright blue. Under ordinary conditions, cuprous salts are less 
stable than cupric salts, but at high temperatures, cuprous salts are 
the more stable. 

2Cu 2 O + O 2 *=* 4CuO 

red black 

Cupric oxide is used in removing sulfur compounds from crude oil 
by a method invented by Frasch. 

How to test for the cupric and cuprous ions. A simple test for 
detecting a cupric salt in a solution is to place an iron nail in the 
solution. A layer of red copper deposited on the nail indicates the 
presence of the divalent copper ion. Why? 

Another test involves the use of ammonia, which immediately 
turns cupric salts deep blue. Cuprous salts in solution may be iden- 
tified also by the addition of ammonia water. Cuprous salts will show 
no color change until the solution is shaken in contact with air. As 
the cuprous ion oxidizes to the cupric ion, the blue color gradually 
grows darker. 

Copper sulfate, the most important salt of copper. Anhydrous 
copper sulfate is white, but when it is crystallized from water solu- 
tion, it unites with five molecules of water and appears as a blue 
crystal.* When heated, this hydrate loses its water of crystalliza- 
tion (water of hydration) , and becomes a white powder. This reac- 
tion is reversible. 

CuSO 4 5H 2 O * CuSO 4 + 5H 2 O 

Copper sulfate is prepared by moving a perforated bucket con- 
taining copper in and out of dilute sulfuric acid. Although dilute 
sulfuric acid has no effect on pure copper, the oxygen of the air re- 
acts with the copper, forming copper oxide, which in tutn reacts 
with the dilute sulfuric acid, forming copper sulfate. 

Minute amounts of copper are almost universally found in living 
cells. Without it man gets anemia, and plants cannot synthesize 
chlorophyll. In large quantities, however, it is poisonous to lower 
forms of life. When copper suifate is added to water, algae and other 
small water plants are killed. Therefore it is used in the purification 
of water. Cloth bags containing copper sulfate often are suspended 
from a boat and pulled through outdoor reservoirs in which drink- 
ing water is stored. 

* The names bluestone and blue vitriol for Crystallized cbppe^ sulfate are now 
considered archaic. 


An excellent fungicide, it is used also (mixed with calcium hy- 
droxide as bordeaux mixture) in spraying fruit trees and potato 
plants. Copper sulfate is used also in refining blister copper and in 
copper-plating (see page 427) . 

How copper is used in printing. The words on a page of this book 
are first set in type. But type metal is so soft that it wears down in 
printing more than 25,000 or 50,000 copies. Therefore, a mold of a 
type page is made in wax. This mold is carefully dusted with a thin 
coating of graphite. The graphite-covered mold is then hung as the 
cathode in an electroplating tank. A plate or bar of pure copper is 
used as the anode and a solution of copper sulfate as the electrolyte. 

A direct current of electricity is turned on until a very thin coat- 
ing of copper has been deposited on the graphite-covered mold. The 
mold with the attached copper shell is removed from the tank and 
the wax is melted away. The thin copper shell is backed with an 
alloy of lead, or type metal, making a plate strong enough to be 
used in printing. This process of using electricity and copper to make 
a printing plate from type is called electrotyping. The finished plate, 
known as an electrotype, can print more than 100,000 copies. 

For long press service or for resistance to abrasive inks, a thin de- 
posit of nickel is made against the graphite-covered wax-mold sur- 
face before the copper shell is formed. Such nickel-faced electro- 
types are known as nickeltypes. Copper or nickel electrotypes may 
be chromium plated for extremely long press runs or severe service. 


Doolittle, Dortha B. "Women in Science." Journal of Chem- 
ical Education, April, 1945, pp. 171-174. Discusses work of Ma- 
ria Mitchell, Ellen H. Richards, Florence Sabin, Caroline 
Herschel, and many others. 

Thorpe, T. E. Essays on Historical Chemistry, pp. 185-205. 
The Macmillan Co., New York, 1923. The life and work of 
Michael Faraday are included. 


Group A 

1. What properties make Cu the most widely used non- 
ferrous metal? 


2. (a) What is the chief use of Cu? (b) Why must Cu for 
most industrial uses be close to 100 percent pure? 

3. (a) What are the three most common alloys of Cu? 
(b) What is the general composition of each? 

4. (a) What f'our areas of the United States produce most 
of our Cu? (b) In the United States where does Cu occur in 
the free state? 

5. (a) What four steps are necessary to obtain Cu from a 
Cu 2 S ore? (b) Write equations indicating the reactions in- 

6. What are (a) matte, (b) blister copper, and (c) copper 

7. (a) Using a diagram and the electron theory, explain 
the electrorefining of Cu. (b) From what four metals is Cu 
completely freed by this process? (c) Why does electrorefining 
of Cu remove these metals? Explain. 

8. Describe the ore-flotation process. 

9. (a) Under what conditions do bronze statues or copper 
roofs get a green coating? (b) What compounds of Cu form 
on their surface? 

10. Why is aluminum ware often preferred to copperware 
for kitchen use? 

11. HC1 has no effect on Cu, whereas HNO 3 dissolves it 
readily. Why? 

12. On the basis of the electron theory, how are the two 
valences of Cu explained? 

13. How would you tell a solution of a cuprous salt from a 
solution of a cupric salt? 

14. What happens to certain cuprous salts when they are 
exposed to air? 

15. (a) What is the difference between, anhydrous CuSO <t 
and its hydrate, crystallized copper sulfate? (b) How can one 
of these salts be changed to the other? 

16. How do you account for the liberation of SO 2 when 
H 2 SO 4 acts on Cu? , t> 


17. How does CuSO 4 help in the purification of water? 

18. (a) What is bordeaux mixture? (b) For \vhat is it used? 


19. Explain the steps in the making of the plates for print- 
ing this book. 

20. Show how the two oxides of Cu illustrate the law of 
multiple proportions. 

21. Cu is used in making silver coins. Why? 

22. What volume of NO is formed by the action of Cu on 
126 g. of nitric acid containing 60 percent HNO 3 ? 

23. What is the percentage of Cu in an ore containing 80 
percent malachite, Cu 2 (OH) 2 CO 3 ? 

24. What simple method is used to reclaim Cu from the 
waste water from copper-mining operations? 

Group B 

25. In plating sheet iron with Ni, why is it necessary to plate 
the Fe first with Cu? 

26. How would you test for the presence of Cu in an alumi- 
num-copper alloy? 

27. Using the electron theory, explain why Cu is so effective 
in reducing HNO H to NO. 

28. Cu is not attacked by dilute H 2 SO 4 , yet CuSO 4 is pre- 
pared commercially by dipping Cu in this acid. Explain. 

29. In the electrolytic refining of Cu, why does the Zn im- 
purity not precipitate on the pure Cu cathode? 

30. Explain the "transmutation" of Fe into Au as "prac- 
ticed" by certain unscrupulous alchemists. 


1. In the market quotations of your daily newspaper, find 
how copper compares in price with steel, Al, Zn, Pb, Sn, Ni, 
and Cr. Make a graph showing price fluctuations of Cu over 
a period of a month. Give reasons for this fluctuation. 

2. Set up a small copper-plating tank, and demonstrate the 
process of copper-plating to your class. Why is this process im- 

3. Name as many household articles as you can that are 
made either of pure Cu or of copper alloys. Is Cu indispensable 
in making such articles? Give reasons for your answers. 

4. Write a two- or three-page report on the life of Michael 
Faraday and his discovery of the principle of the dynamo. How 
has this discovery affected our way of living? 



. . . Iron breathes the air f burns it- 
self up in oxygen and so gives its 
life that we may live. John Ruskin, 

. . . Rust and corrosion mean an 
enormous loss to Americans, greater 
than that caused by fire and flood 
combined, a loss of at least several 
billion dollars a year. Rust is a skin 
disease. Corrosion is an infectious 
internal disease like tuberculosis. 
Iron Trade. 

Have all the elements been discovered? Seventy of the 92 natural 
elements are metals. Between 1923 and 1945, the six elements neces- 
sary to complete Mendeleyeff's periodic table were discovered. Of 
these, five turned out to be metals. Two Italian investigators, C. Per- 
rier and Emilio Segre, obtained element 43, technetium, in 1937. This 
was the first synthetic element ever created. In 1940, element 85, 
astatine, was produced in the cyclotron, and during the research on 
nuclear fission, 1940-1954, eight new elements (Nos. 93-100) were 
produced in the United States (see page 187) . Three of the most 
widely used metals iron, copper, and aluminum have already 
been discussed. In this chapter we shall not consider most of the 
remaining metals, many of which are extremely rare and from our 
point of view relatively unimportant, but shall focus our attention 
upon those whose abundance, properties, and uses make them more 
or less familiar. 

How science fights corrosion of metals. Rust and corrosion are two 
great destroyers of property. As soon as most metals are exposed to 
air, water, or chemicals, an attack is launched upon them. Rust at- 
tacks their surfaces, and corrosion eats its Waty into most metals. 

Science battles corrosion and rust with various weapons, the most 
important of which are: (1) putting a coating of a^cnore resistant 




metal such as zinc, tin, chromium, nickel, silver, cadmium, and 
gold, over a less resistant one; (2) alloying a metal to form a product 
that is more resistant; (3) covering a metal with a thin film of its 
oxide or with a coating of paint, lacquer, or enamel; (4) adding 
some organic inhibitor such as sodium chrome glucosate to the radi- 
ator of a car to prevent corrosion of the cooling system; and (5) sac- 
rificing one metal to protect another in cathodic protection. The 
active metal magnesium protects the less active metal iron from cor- 
rosion due to small, local electric currents. Buried iron pipelines 
are protected from corrosion by attaching to the iron pipes blocks 
of magnesium metal by means of an insulated copper cable (see 
Fig. 97) . Magnesium anodes or an impressed electric current is used 
to protect canal gates, water tanks, and the metal bottoms of our 
moth-ball fleet lying in several waters. 

A powerful film to protect sheet iron. Tin is a soft, white metal 
that melts at a comparatively low temperature (232 C.) but, com- 
pared with other metals, it is highly resistant to oxidation by air 
and to acids found in certain kinds of foods, such as fruit and to- 
matoes. Furthermore, tin is scarcely affected by dilute hydrochloric 
or sulfuric acids. 

Because of its chemical inactivity, tin is used extensively in cov- 
ering light-gauge sheet steel to protect it from corrosion. First, the 
steel is scoured with sand, then pickled in dilute sulfuric acid, and 
washed thoroughly to remove the acid. It is then dipped in a bath 
of molten tin, on which is a layer of oil that prevents any oxidation 
of the steel during the tin-plating process. Most of this tin plate is 
used in the manufacture of tin cans for food canning and the oil 
industry. The weight of tin in a tin can produced by this hot dip 
process is hardly more than one percent of the weight of the can. 

Tin plate is also produced electrolytically by a new process devel- 
oped as a result of the tin shortage during World War II. This proc- 
ess is much faster than the older hot clip process and deposits the 

Fig. 97. Cathodic protection of buried iron pipe. 

Mg of 


'soil and 

fn - 





tin in thinner, more even coatings, thus saving more than one mil- 
lion pounds of tin a year. In this process, tin anodes are used to sup- 
ply the tin for plating the cold-rolled steel strip that forms the cath- 
ode. As much as 1000 feet a minute can be coated with a film only 
30-millionths of an inch thick. 

Tin is used also in making bronze, solder, Babbitt metal, and 
pewter, an alloy of 75 percent tin and 25 percent lead used com- 
monly during Colonial days for tableware. The comparatively high 
cost of tin prevents its more extensive use. Tin salts are used for 
weighting silk. 

Little or no tin is mined in the United States; we are dependent 
for our entire tin supply on mines in Bolivia, Indonesia, and British 

Tin has two valences. The stannous ion has a valence of two and 
the stannic a valence of four. (Consult Table 14 for other facts about 
the metals discussed in this chapter.) 

Nickel and nickel-plating. Nickel, a silvery-white metal, is not 
oxidized by dry air and is only slowly attacked by moist air. It is 
therefore used extensively for plating. Nickel resembles tin in that 
it is scarcely affected by hydrochloric or sulfuric acids. Nickel-plat- 
ing is done electrolytically, using a nickel anode, and a nickel salt 
as electrolyte. The metal is used also as a catalyst in the hydrogena- 
tion of oils, in the nickel-alkali storage battery (Edison cell) , and in 
the cadmium-nickel alkaline storage battery. 

Nickel alloys and cobalt are important. "Monel metal," which con- 
tains 70 percent nickel, 28 percent copper, and two percent iron, is a 

strong, acid-resisting alloy used extensively. It is only one of the 
many important nickel allies on the nonrust battlefield. Our own 
five-cent piece contains 25 percent nickel alloyed with copper. Coins 
used in more than 20 countries are pure nickel. 

Many steel alloys contain nickel. "Invar," which expands very lit- 
tle on heating and is therefore used for lead-in wires in electric-light 
bulbs, and in instruments that measure accurately, is a steel alloy 
containing 35 percent nickel. "Permalloy" and "Hipernik" are steel 
alloys containing high percentages of nickel. Because these alloys 
are more magnetic than iron, they are used in transformers. "Alnico," 
a family of nickel alloys that contain nickel, iron, cobalt, and alumi- 
num, can lift 500 times their own weight of iron. Alnico magnets are 
used in radio and TV receivers, and in magnetic tape recorders. 

An even more powerful magnet is made of platinum and cobalt, a 
metal closely related to nickel. Cobalt is a very strategic metal used 
in many important alloys. The stellites made of cobalt, chromium, 
and wolfram are used for high-temperature, high-speed cutting tools. 
Vitallium containing cobalt, chromium, and molybdenum is used in 
jet engines and gas turbines which are subjected to extremely high 

Chromium, the tarnishproof plating. Chromium is used in a num- 
ber of very important steel alloys. Its extreme hardness, high luster, 
and very high resistance to corrosion, have brought chromium to the 
forefront as a plating metal. Automobile headlight shells, bathroom 
fixtures, and tarnishproof jewelry, are only a few of the articles now 
plated with chromium rather than with nickel. The electrolyte used 


Use by approximate 


Use by approximate percentage 

in this plating is chromic acid containing a catalytic agent. Nichrome 
is an alloy of nickel, iron, manganese, and 1 1 percent chromium. 

Chromium is obtained by reducing chromite, Cr 2 O 3 FeO, with 

Cr 2 O 3 4- 2A1 - A1 2 O 3 + 2Cr (see thermit process) 

or by electron reduction. Though neither fuming nitric acid nor 
even aqua regia attacks chromium, warm dilute sulfuric acid or hy- 
drochloric acid reacts with it, liberating hydrogen. 

2HC1 + Cr -> CrCl 2 + H 2 T 

Chromium has several different valences, ranging from two to six. 
Most chromium salts are highly colored (chroma means color) . 

Zinc and galvanized iron. One of the oldest methods of protecting 
iron from rust and corrosion is to coat it with zinc. This is done in 
one of three ways: (1) immersing sheet iron that has been pickled 
and then thoroughly cleaned in a bath of molten zinc; (2) deposit- 
ing zinc on the iron electrolytically; (3) packing the article to be 
coated in zinc dust and heating to about 300C., a process called 
sherardizing. The zinc condenses onto the metal. 

Although zinc is higher in the replacement series than iron, it 
acts as a protective coating, because the thin layer of zinc oxide or 
zinc carbonate that forms on exposure to air sticks to the metal. 




This coating prevents further corrosive action, thus protecting the 
easily corroded iron. Iron coated with zinc is the familiar galvanized 
iron used in making pails, cans, and pipes. The cathode of all dry 
cells is made of zinc. Large quantities of zinc are used also in the 
alloys bronze, brass, and nickel (German) silver. 

Zinc dust catches fire when heated and burns with a bright light, 
forming zinc oxide, ZnO. This oxide, under the name zinc white, 
is used in making paints and in the manufacture of automobile tires 
its high heat-conductivity helps to keep the tire cooler and thus 
prolongs its life. It is also used in zinc ointments, face powders, deo- 
dorants, and similar cosmetics where it furnishes covering and anti- 
acid and adhesive properties. 

How zinc is freed from its ores. Zinc is not found free. Its chief 
ore is zinc blende, ZnS. Its metallurgy is simple. After first concen- 
trating the ore by flotation, the ZnS is roasted to convert it into the 
oxide. This oxide is then mixed with finely powdered coal and 
heated in clay retorts. This process reduces the zinc oxide to zinc. 
The zinc vaporizes, and is then condensed into a solid, melted, 
poured into ingots, and sold as spelter. 

2ZnS + 3O 2 - 
ZnO + C 

2ZnO -f 2SO 2 T 

Zn + CO T 

Natural gas, instead of coal, is also used as the reducing agent. 
The advantages of this newer process are: (1) natural gas is found 


nearer zinc mines than coal; (2) a lower temperature is sufficient to 
reduce the ZnO; (3) the process is continuous rather than intermit- 
tent. The chief chemical reaction is: 

ZnO + CH 4 -> Zn + CO | + 2H 2 1 

Cadmium resembles zinc. Many zinc ores contain from one to 
two percent of a metal that closely resembles zinc in appearance and 
properties. This metal, cadmium (Cd) , is freed from its ores in the 
same manner as zinc. Cadmium is somewhat harder than zinc and 
is superior to zinc as a plating metal, because the finished plate is 
free from the pinholes so frequently found in zinc plate. Why do 
pinholes lessen the effectiveness of zinc-plating? 

The chief use of cadmium is in electroplating. Coatings of cad- 
mium one-third as thick as zinc provide as much protection. It is 
used also in making the yellow pigment, CdS, for paints and printing 
inks. Alloys of cadmium are used extensively in making bearings and 
for other uses in which a relatively low melting point is desired. 

Silver, the noble white metal. Silver, which is very low in the 
replacement series, occurs free (usually alloyed with gold) as well as 
in a variety of ores, chief of which is the sulfide of silver, argentite. 
This ore is roasted, and the oxide is reduced with charcoal. 

When silver is found as part of a lead sulfide ore, it is separated 
from the lead by the Parkes process, in which the ore is roasted and 
then reduced to a silver-lead alloy. To this molten alloy about two 
percent of zinc is added. Silver, being very much more soluble in 
zinc than in lead at the temperature of the molten alloy, mixes with 
the zinc and rises to the top of the melt. The silver-zinc mixture is 
then skimme.d off and heated, and the zinc is distilled off. 

Silver is called the white metal because it can be obtained as tiny 
white crystals. It is the best-known conductor of heat and electricity 
and, next to gold, the most malleable and ductile of all metals. 
Oxygen has no effect upon it, but ozone oxidizes it. Why? 

6Ag + 3 -3Ag 2 

Removing tarnish from silverware. The tarnishing of silver ex- 
posed to air is caused by the formation of a film of brownish-black 
silver sulfide, Ag 2 S, formed by the action of the sulfur or sulfur com- 
pounds in air or in food (yolk of eggs, mustard) . This tarnish may 
be removed by heating the tarnished silverware in a water solution 
of table salt, and washing soda, Na 2 CO 8 , in contact with aluminum. 
Because an aluminum pan is badly stained by this process, it is 


preferable to use either a piece of aluminum in an enameled pan or 
else a cheap aluminum pan. The hydrogen that is liberated reduces 
the silver sulfide to pure silver. The cleansing action may be repre- 
sented by the following chemical equations: 

3Na 2 CO 3 + 3H 2 O + 2A1 -> 2Na 3 AlO 3 + 3CO 2 + 3H 2 

H 2 + Ag 2 S -> 2Ag + H 2 S T 

Flannel cloth impregnated with silver phosphate protects silverware 
wrapped in it. H 2 S attacks the silver phosphate instead of the silver- 
ware. Silver behaves like copper in the presence of the common acids 
(see page 431). 

Silver in jewelry and coinage. Because of its high cost and pleasing 
appearance, silver is used in jewelry and coinage. Because it is too 
soft to stand hard usage, it is alloyed with copper to make more 
durable coins. The amount of silver in alloys is stated in terms of 
fineness, which indicates the number of parts of silver in 1000 parts 
of the alloy. American silver coinage is 900 fine; that is, it contains 
90 percent silver and ten percent copper. Sterling silver is 925 fine, 
containing only 7.5 percent copper. 

At one time, China and India required much of the exported 
silver for use as a monetary standard. What is the present price of 
silver in the United States? 

Silver-plating. The making of silver-plated articles, especially table- 
ware, consumes much of the silver produced annually. The process 
of silver-plating is similar to copper-plating. The article to be plated 
is hung from the cathode. A sheet or bar of pure silver is used as the 
anode. The electrolyte is a solution of silver potassium cyanide 
which, although extremely poisonous, is used instead of other elec- 
trolytes because it gives the most even and durable coating (see il- 
lustration on page 444) . The manufacture of silver chloride and 
silver bromide, used in photography, required more than 400 tons 
of silver in a recent year. 

Silver is now employed more widely by industry than by the arts 
in the United States. Silver alloy brazing of airplane, radar, and radio 
parts is responsible for this shift. 

Gold, the yellow metal that resists corrosion. Gold, which is found 
in the free state, was known to the earliest civilizations, as evidenced 
by the golden ornaments of elaborate workmanship which they have 
left to us. The gold fields of Nubia and of Asia Minor were the 
sources of the riches of Croesus, wealthy king of ancient Lydia in 
Asia Minor, who lived about 2500 years ago. Gold is obtained in var- 
ious parts of the world from quartz veins which must be crushed, 

K+Ag+fCN) 2 - solution 

Object to 
be plated 

Fig. 98. Electrolytic process of 
silver-plating. How does this 
compare with the copper- 
plating process? 

and from river sands or gravels. Today the greatest center of gold 
production is the Transvaal in Africa. 

When gold occurs in river sands or gravels, the mining operation, 
called placer mining, is simple. About three-fourths of the world's 
gold is obtained by placer mining. The gold prospector tests his 
"dirt" by panning, that is, stirring the sand with water in a pan to 
enable the heavy gold to settle to the bottom. A single pan in the 
Klondike sometimes yielded a pound of pure gold. The largest gold 
nugget ever found was picked up in Victoria, Australia, in a wagon 
rut only a few inches below the surface of the ground. It was called 
Welcome Stranger, and weighed 157 pounds. 

Gold ores usually contain silver also. The two are separated by 
adding nitric acid in which only the silver is soluble. Gold is often 
refined electrolytically (see pages 427-428) , the electrolyte being a 
solution of gold chloride containing some hydrochloric acid. 

Gold is not attacked by air, water, or any one of the common acids. 
Gold dissolves in aqua regia, going into solution as gold chloride. 
Alkali cyanides, such as highly poisonous sodium cyanide, NaCN, 
also dissolve it. About 60 percent of the gold in gold ore is extracted 
by mercury, with which gold forms an amalgam (page 446) . NaCN 
is used to remove the remaining gold from the ore. 

Gold in jewelry and coinage. Like silver, gold is used in jewelry 
and coinage. Because of its softness it is alloyed with copper to pro- 
duce greater hardness. Gold coins formerly used in the United States 
contained ten percent copper. For 100 years the American gold dol- 
lar was backed by 23.22 grains of pure gold (1 troy ounce = 480 
grains) . Franklin D. Roosevelt in 1934, under powers conferred upon 
him by the Congress, changed the gold content of the dollar to 13.71 
grains, worth slightly more than 59 cehts of the former dollar. The 
price of gold was changed from $20.67 per troy ounce to $35. 

The purity of gold is measured in carats. Pure gold is 24 carat. 
Eighteen-carat gold, used extensively in jewelry is 18/24, or 75 per- 
cent gold. About one-half of the gold mined each yeaY is stored to 
back up currency; the rest is used in the arts and crafts. 



Gold-plating and other uses of gold. Much jewelry and many 
other small objects are made of one of the cheaper metals plated 
with gold. Gold-plating on an object prevents tarnishing and makes 
possible a brilliant polish. Gold plating is carried out in the same 
way as silver-plating. The anode is pure gold and the electrolyte is 
gold potassium cyanide. Liquid gold, used to decorate pottery and 
earthenware, is an organic salt of gold dissolved in various oils and 
applied to the surface of the glaze. After firing, a brilliant film of 
metallic gold is left. Gold leaf is used also for decorative purposes, 
for lettering bookbindings, letters on windows, and the like. White 
gold is an alloy of gold and some other metal such as palladium or 
nickel. Much gold is used in the production of electric equipment 
and in color photography. 

Mercury, the only liquid metal. Occurring in small amounts in the 
free state but obtained easily from its ores, the metal mercury was 
known to the ancients. Cinnabar, HgS, a brownish-red compound, 
is the chief ore of mercury. The chemistry of its extraction is simple. 
The ore is heated to vaporize the mercury. The vapor is then con- 
densed, purified, and shipped in flasks of 76-pound capacity. 

So dense is mercury (density = 13.6 grams per mil HI her) that 
iron floats on it. Mercury remains unchanged in dry air, oxygen, and 
carbon dioxide. 

As you know, when mercury is heated, it changes to red mercuric 
oxide, HgO, which upon further heating decomposes back into its 
elements. It was this simple chemical change that led Lavoisier to 
the true explanation of burning. Dilute nitric acid readily attacks 
mercury, and, in general, mercury reacts with the acids much as 
copper does. 

3Hg + 8HNO 3 > 3Hg(NO 3 ) 2 + 2NO | + 4H 2 O 

Westinghouse Electric Corporation 

A research scientist placing pure gold 
wire around the valve opening of a 
vacuum tank used in the production 
of zirconium. Gold was the only mate- 
rial found satisfactory in making an 
air-tight, high-temperature, corrosion- 
resistant seal. 





Mercury combines directly with sulfur and the halogens. It is 
sometimes mixed with metals forming mercury alloys, called amal- 
gams. One such amalgam contains silver, and is used by dentists to 
fill teeth. 

Adding one valence electron produces a poison. The element 
mercury has two valences and hence forms two series of salts. Mer- 
curous chloride, HgCl, is a white crystalline salt, calomel, used in 
medicine. Mercuric chloride, HgCl 2 , also is a white crystalline com- 
pound. It is a violent poison and is known as bichloride of mercury, 
or corrosive sublimate. It is such a powerful germicide that it de- 
stroys all known bacteria. The Hg++ ion is poisonous, but the Hg f 
ion is not. Mercury compounds, formerly used in hatmaking, pro- 
duced mental disturbances. Hence the term "mad as a hatter." "Mer- 
curochrome" is an antiseptic consisting of a complex organic com- 
pound containing 23 percent mercury. 

From measuring instruments to turbines. Mercury, also known 
as quicksilver, is commonly used in thermometers, barometers, and 
other scientific instruments. It is used in preparing the mercury salts 
mentioned above as well as in making mercury fulminate, Hg 
(CNO) 2 , an unstable compound used as a detonator for explosives 
(see page 270) . 

Mercury may be used in boilers and turbines. A huge electric gen- 
erator using mercury vapor instead of steam, is in operation in Sche- 
nectady, New York. The mercury vaporizes at high temperatures 

and its vapor drives a turbine. The exhaust mercury vapor passes 
through a steam boiler in which its high temperature produces su- 
perheated steam, which, in turn, drives a steam turbine. The system 
is closed and the process is continuous; mercury is condensed, re- 
turns to the mercury boiler, and is again vaporized. 

Sunlight and fluorescent lamps. Other uses of mercury are in sun- 
light and fluorescent lamps. In the sunlight lamp an arc is formed 
between two wolfram electrodes in the bulb, and the pool of mer- 
cury that is vaporized gives off light rich in ultraviolet radiations, 
which have health-promoting value. Radiations from such lamps 
also kill airborne and surface bacteria. 

Mercury is used in the mercury vapor lamp employed in photog- 
raphy and in mercury vapor vacuum tubes that are used as rectifiers 
to convert alternating current (ac) to direct current (dc). Such tubes 
are used in storage-battery chargers, in a-c radios, and in other in- 
stallations requiring the conversion of alternating current to direct 

The fluorescent lamp is one form of "electric discharge" light 
source. It consists of a tubular bulb with an electrode sealed in each 
end. A flow of electricity takes place through a drop of mercury that 
has been vaporized. The ultraviolet light produced causes certain 
powders, called phosphors, that adhere to the walls of the tube to 
fluoresce, that is, to emit visible light while absorbing the invisible 
ultraviolet radiations. Almost any color or tint can be produced by 


Use by approximate 

Electrolytic preparation 
of chemicals 

Industrial and control 

Electrical equipment 

Antifouling paint 

General laboratory 

Dental preparations 


mixing these phosphors, such as zinc silicate or magnesium tungstate. 
Phosphors are also used to coat TV, radar, and x-ray screens. A mov- 
ing beam of electrons excite and light up the phosphors. 

Fluorescent lamps are much more efficient than hot-filament lamps. 
The hot-filament, or incandescent, lamp is really more a heater than 
a light source. Light is produced only by heating the filament to 
incandescence. More than 90 percent of the electricity supplied is 
consumed in heating the filament and less than ten percent is used 
in producing light. Fluorescent tubes are more than three times as 
efficient as incandescent lamps; that is, they give more than three 
times as much light for the electricity consumed. 

Lead and manufacture of white lead. Lead, a slightly lustrous, 
bluish-white, soft and malleable metal, occurs free very.rarely. When 
exposed to air, it is quickly coated with a protective layer of lead 
oxide or lead sulfide. Lead reacts slowly with most corrosive acids. 

Excessive amounts of lead salts form cumulative poisons that may 
affect the entire human body, especially the nervous system. The 
dangers of lead poisoning were known for a long time. Wedgwood, 
the famous maker of glazed pottery and friend of Priestley, wrote in 
1775, "I will try in earnest to make a glaze without lead/' and he did. 
Lead poisoning, once a serious occupational disease, is now less com- 
mon, thanks to mechanical safeguards and humanitarian legislation. 

The principal uses of lead, listed in order of the quantities con- 
sumed, are storage batteries, cable covering, tetraethyl lead for gaso- 
line, and pigments (litharge and white and red lead) . Lead is also 
used for pipes, tank linings, nuclear reaction and x-ray protective 
plates, shot, reaction chambers for the manufacture of sulfuric acid, 
and in alloys such as solder, type metal, Babbit metal and terne metal 
(sheet iron; covered with a layer of lead-tin alloy) . 

The manufacture of white lead was carried on as early as 300 B.C. 
From an old manuscript of this time, the following description is 
given: "Lead is placed in earthen vessels over vinegar, and after it 
has acquired a rust, they open the vessel and scrape it off. Then they 
beat it to a powder and boil it." This is essentially the method used 
in the Old Dutch process. The chemical reactions involved are: 

1) Lead reacts with vinegar (acetic acid) , forming lead acetate. 

2) Lead acetate reacts with carbonic acid, formed by the fermen- 
tation of tanbark which is placec^over the earthenware pots, and 
white lead is formed. The average composition of white lead is 
Pb (OH) , - 2PbC0 8 . 

Modern production methods, such as the Carter and Sperry proc- 
esses, are improvements on the Old Dutch process. These newer 

Lead Industries Association 

The television tube shown here being welded contains as much as 3O percent 
lead by weight. 

methods have cut clown the time required in the manufacture of 
white lead from a few months to a few hours. 

Hiding is desirable in paint. Paint is essentially a mixture of body 
pigment and a liquid vehicle. To these may be added a thinner to 
reduce the mixture to painting consistency, a drier to accelerate the 
drying of the vehicle, and a tinting, or coloring, pigment. 

Linseed oil, obtained from flaxseed, is the most commonly used 
vehicle, or medium. During the drying process, it polymerizes with 
the aid of the oxygen of the air, forming an elastic, impervious, horn- 
like film. Tung oil, from the Chinese tung tree, and soybean oil are 

Non-corrosive automobile gasoline tanks, made from terne-plate, a lead-tin alloy. 

Lead Industries Association 





often used instead of linseed oil. Turpentine is the most commonly 
used thinner. It evaporates after the paint has been applied. Among 
the most frequently used driers are certain oxides, such as MnO 2 and 
Pb,O,, and certain metallic soaps, all of which act catalytically to 
hasten the drying process. Colored pigments include such chemical 
compounds as yellow cadmium sulfide, CdS; chrome yellow, PbCrO 4 ; 
Venetian red, Fe L >O :j ; carbon black; Prussian blue; and red lead, 
Pb 3 O 4 , which is the standard protective paint for iron and steel. 

A frequently used body pigment, or base, is white lead. This 
opaque compound is ground with linseed oil and the required 
amount of drier, color, and thinner. A filler, or extender, is an in- 
ert material added to paint to give greater bulk and decrease cost. 
Common fillers are CaSO.^ BaSO 4 , kaolin, fine white sand, and CaCO 3 . 
However, when used in excessive amounts, fillers impair the life of 
the paint film. 

In the production of paints, one of the chief objectives is the 
achievement of hiding power, or the ability to cover completely the 
underlying surface. For this purpose titanium dioxide, TiO 2 , a white 
pigment with a high whitening power, is added to the paint mixture. 

The element titanium occurs abundantly in the United States in 
the ore ilmenite, FeTiO ... Titanium dioxide is produced from ilmen- 
ite, without the production of titanium as a metal, by several chem- 
ical reactions ending in large rotary kilns. In addition to being used 
in paints, titanium dioxide is also used extensively in the produc- 
tion of paper products of high opacity, in making linoleums, rub- 
ber articles, cosmetics, glass, and so forth. 

Lead and storage batteries. The lead storage cell is a device for 
changing chemical energy into electric energy. In a six-volt storage 

battery, there are three cells. It contains a lead plate with pockets 
of spongy lead as the negative electrode, and a lead plate with pock- 
ets of lead peroxide, PbCX, as the positive electrode. 

Both plates are immersed in a dilute solution of sulfuric acid 
(sp. gr. 1.2) . When the exterior terminals of the plates are connected 
by a copper wire, the chemical changes that take place cause a transfer 
of electrons that produces an electric current. The spongy lead gives 
up electrons which travel to the PbCX. We say that the current moves 
from the plus pole to the minus pole. Actually, however, the flow of 
electrons is from the negative to the positive pole. As the battery is 
used, both plates become covered with PbSO 4 . The changes that 
occur during the discharge of the cell are: 

Negative plates: Pb + H 2 SO 4 -> PbSO 4 + 2H+ + 

Positive plates: PbO 2 + H 2 SO 4 + 2H+ + 2e -> PbSO 4 + 2H 2 O 


As you can see, each of these reactions removes some H 2 S() 4 which 
changes to H 2 O, and the electrolyte, therefore, becomes more and 
more dilute, and its specific gravity falls. Hence, we can test the 
strength of the cell by testing the specific gravity of the solution by 
means of a hydrometer (see Fig. 100) . When the cell has run down, 
as indicated by the dilution of the electrolyte, it can be recharged 
by passing a direct electric current through it in the opposite direc- 
tion. This reverses the reactions given above. 

A storage battery should not be discharged completely, nor charged 
or discharged too rapidly. Why? Water must be added to replace the 
water lost through gradual evaporation, and chemical action. 

Bismuth helps fight fire. Bismuth is a white, lustrous, brittle metal 
with slightly reddish tinge. It occurs both free and combined. It is 



Lead plate with pockets 
of spongy lead 

Lead plate with pockets 
containing PbO 2 


Fully discharged 1150 
Half charged 1225 
Fully charged 121 

Acid level 

now shown 

at 1280 

Fig. 99. The lead storage cell. 

Fig. 100. Testing a storage battery. 

used chiefly in making low-melting-point alloys such as Wood's metal, 
which contains 50 percent bismuth and melts at 60C. Electric fuses, 
safety plugs in boilers, and other automatic devices for protection 
against fire or explosion contain this metal. If a fire breaks out in 
a building equipped with automatic water-sprinklers, a temperature 
of 60 C. is reached quickly. The bismuth alloy in the plug then 
melts, setting a stream of water loose and at the same time breaking 
an electric circuit, which sends out a fire alarm. 

Bismuth subnitrate, Bi (OH) 2 NO 3 , is used for digestive disturb- 
ances, and as a suspension for x-ray of the digestive tract because 
it is opaque to x-rays. Cosmetics containing bismuth compounds may 
be dangerous since the compound, bismuth subnitrate, hydrolyzes in 
contact with perspiration, forming nitric acid. A fungicide spray com- 
posed of a bismuth compound is effective against mildew diseases 
ol tobacco, potato, and other plants. 

Bismuth, antimony, and arsenic occur in the periodic table (see 
page 162) in the same group as phosphorus and nitrogen. Bismuth 
is a true metal; phosphorus and nitrogen are typical nonmetals. 
Arsenic and antimony, though resembling metals in certain physical 
properties, are very close to the nonmetals in other properties. They 
are borderland elements, acting under certain conditions as metals 
and under others as nonmetals. 

Antimony, a metal that expands on solidifying. Antimony is sil- 
very, brittle, and crystalline and, quite unlike most metals, expands 
when it solidifies. Its chief use is in alloys such as type metal and Bab- 
bitt metal. 

Type metal, used in printing, is 'made of lead and a small amount 
of tin and antimony. When molten type .$&$} solidifies, it expands 
slightly, thus spreading into the sharp corners of the mold, producing 
a sharply defined cast. Most metals shrink when solidified and can- 
not be cast but must be stamped or struck out ofi die. 




Babbitt metal, named after its American inventor, Isaac Babbitt, 
contains antimony and tin and a small amount of copper. It is used 
in making anti-friction bearings for machinery. Compounds of anti- 
mony are used to make canvas and other textiles rotproof and fire- 
proof. Antimony preparations were used as medicinal agents back in 
the Middle Ages. 

Arsenic, a poison and a preservative. Arsenic is a steel-gray, brittle, 
metallic-looking solid that is found both free and combined. Like 
bismuth, it is seldom used as the pure element. Lead shot is made 
harder by the presence of one percent of this metal. When heated in 
air, it forms arsenic trioxide, exhibiting, like bismuth and antimony, 
a valence of three toward oxygen. 

4As -f 3O 2 > 2As 2 O 3 (white arsenic) 

White arsenic is a powder that is used in glassinaking and as a 
preservative in the mounting of skins. The poisonous nature of the 
compounds of arsenic is used in man's battle against the insect pests 
which, by attacking his cattle and crops, cause losses of billions of 
dollars yearly. Insects destroy approximately ten percent of man's 
food and fiber crops each year. Suspensions of lead arsenate and cal- 
cium arsenate are widely used agricultural insecticides for spraying. 
Paris green is an arsenic salt of copper. 

Arsenic compounds are stomach poisons and are used against in- 
sects that eat foliage. Contact insecticides such as lime-sulfur and 
nicotine dust are employed against sucking insects. The great advan- 
tage of calcium arsenate is its low price which makes possible its use 
in cotton fields against the destructive boll weevil. It also kills the 

- ... ' . " ' 7 " 5tf- rs v " v , / ;;*- - **'<*-, r / ^ ' 




tomato worm. Lead arsenate is effective against insects that attack 
fruits, flowers, and shrubs. Fruits and vegetables exposed to excessive 
amounts of such poisonous sprays should be carefully and thoroughly 
washed before being eaten raw or cooked. 

Platinum, one of the very heavy metals. Antonio de Ulloa (da- 
dbl-yo'a) , 200 years ago, described a strange metal that he had seen 
in Peru. Because its value was not appreciated then, it was used 
to adulterate gold. When a heavy gray nugget of this metal was 
brought to Europe and studied carefully, it was found to be a new 
element, platinum. 

The properties of platinum make it very valuable. It is very resist- 
ant to corrosion. Like gold, platinum is attacked only by aqua regia, 
which dissolves it, forming platinum chloride. Platinum crucibles 
and dishes should not be used for melting alkalies or metals of low 
melting point, since platinum is attacked by these substances. Be- 
cause of its high resistance to the passage of electricity and its high 
melting point (1750C.) , platinum is used in making electric fur- 
naces, resistance thermometers, spark-plug tips, rayon spinnerets, and 
x-ray and radio tubes. 

One of platinum's most important uses is as a catalyst in the con- 
tact process for manufacturing sulfuric acid, and in making nitric 
acid by the oxidation of ammonia. Because of the rarity, beauty, and 
durability of platinum, it is widely used in making jewelry. Surgeons 
use it to replace parts of bones, and dentists use it in bridgework. 
Platinum salts are used to make photographic prints more permanent 
and beautiful than those made by the use of silver salts. 

Platinum has a density of 21.45 grams per cubic centimeter. To- 
gether with the other members of the heavy metal family (iridium, 
osmium, ruthenium, rhodium, and palladium) , it is mined in the 

Westinghowc Electric Corporation 

This hard, gray, coke-like lump is 
titanium sponge, the mid-point in 
the production of commercially 
useful titanium metal from rutile, 
ilmenite, or other ores. 


Us* by 

gravels of river beds, chiefly in the Ural Mountains in the Soviet 
Union and in Canada. Some platinum is mined in Brazil, and in 
this country in California and Oregon. 

In the hope o( : freeing itself from the necessity of using metals, 
such as platinum, which are both expensive and subject to wide 
fluctuations of price, industry has turned to new alloys that are 
cheaper to produce. This technological war among the metals has 
multiplied the kinds of materials available for man's use and has 
ushered in the age of alloys. In this war of the metals, less well-known 
metals have come into wider use. 

Many "rare" metals are really very abundant. Titanium, zir- 
conium, vanadium, wolfram, and tantalum are actually more abun- 
dant than lead in the earth's crust. Zirconium is more than twice as 
abundant as copper, and thirteen times as plentiful as lead. It is 
highly resistant to many acids and is used in our submarine nuclear 
reactors because it is a low neutron absorber. Some new casting alloys 
contain this metal. Titanium, fourth most abundant metal and new- 
est ol the structural metals, is about half as heavy as steel, resists sea- 
water corrosion better than steel, and has great tensile strength. Like 
zirconium, it is highly reactive at high temperatures, becoming 
brittle and hard to handle. Gallium, 150 times as abundant as silver, 
melts at body temperature and expands on solidification. 

One of the reasons these so-called rare metals are still not widely 
known has been the difficulty of extracting them cheaply from their 
ores. However, as one recent writer put it, "It is not improbable that 
many of the present well-known metals may, within the next half 
century, become outmoded in favor of certain of the relatively more 
abundant but less-known metals of today." 

Wolfram, the metal used in electric-light bulbs. Wolfram's origi- 
nal name was tungsten, a misnomer since the element is a metal. This 
steel-gray, heavy metal, three times as hard as platinum and with a 




melting point (3300 C.) almost twice as high as that of platinum, 
has gradually replaced platinum in the electric industry. Today, 
wolfram filaments in electric-light bulbs are in general use. 

The United States consumed more than 12,000 tons of this metal 
in a recent year, most of it for the manufacture of high-speed tool 
steel alloys and machinery. Wolfram steel is not only hard but stays 
hard even when red hot. The speed of boring and cutting machines 
was previously limited by our inability to keep the cutting tool cold. 
As the tools became heated, they lost their keen edges. This metal has 
also been used to make wolfram carbide, which, when welded to- 
gether in the form of millions of minute particles (cobalt being the 
welding agent) , makes an excellent abrasive known as "Carboloy." 

Molybdenum, tantalum, germanium, and beryllium. Molybde- 
num (Mo) is similar to wolfram in physical properties, but its melt- 
ing point is lower (2500 C.) , and its lower cost has made it an ac- 
tive competitor of wolfram in the electric field, and in the making of 
high-speed tool steel alloys. Bartlett Mountain in Colorado holds 
about 95 percent of all the world's known store of molybdenum. 

Tantalum (Ta) is another recent addition in the electric field. It 
is a corrosion-resisting metal resembling lead. Its price is only one- 
fifteenth that of platinum, and already uses have been found for it 
in electron tubes, neon tubes, surgical and dental instruments, acid- 
proof coils, electric contacts, art metal, and in the spinnerets used in 
synthetic textile manufacture. 

Germanium crystals are used in tiny transistors which have recently 
replaced vacuum tubes for many uses in electronics. 

Beryllium (Be) , the fourth lightest element, is another newcomer 
that has found its greatest use in copper alloys. These alloys are used 
in making certain tools and springs. Beryllium alloy springs retain 
their resiliency remarkably well. 


t>y opproxii 



Fig. 101. Demonstration of the 
principle of oil-retaining bear- 
ings. When the aspirator bulb 
is squeezed, oil wells up in the 
pores of the bearing to form a 
protective coating. 

Powder metallurgy. Most metal products are made by casting the 
molten metal in molds, by machining a block of the metal to the de- 
sired shape, or by a combination of casting and machining. But metal 
products in great quantity are being made by a newer process that 
consists of (1) pressing metal powder or a mixture of metal powders 
into form, and (2) baking this compressed metal powder form in a 
sintering furnace. Metal products so produced are porous but of great 
strength. Wolfram filaments for electric lamps and electron tubes 
are made by sintering powdered wolfram and hammering the re- 
sulting product into very fine but porous wires. Wolfram carbide 
cutting tools are also made by this process. 

Among the many interesting metal products made by sintering are 
porous bearings that can be filled with oil and yield enough lubrica- 
tion for the life of the equipment. Squeeze the bearing, and oil wells 
up from its pores. Release the pressure and the oil flows back into 
the bearing. Such bearings are called oil-retaining bearings and gen- 
erally are pressed from tin, copper, and graphite. Under great strain, 
they provide the greatest quantity of lubrication; under slight strain, 
very little. The oil in the bearings is not affected by extremes of tem- 

An operator filling a "die 
well' 1 with metal powder. 
When the well is filled with 
powder, the upper punch 
will move down, compress- 
ing the powder into a 
cylindrical bearing such as 
that shown. 

Westinghouse Electric Corporation Bell Telephone Laboratories 

(left) Hairpin-shaped bars of 99.9 percent pure zirconium emerge from the vacuum 
tank in which they have been produced, (right) A scientist studies x-ray plates to 
determine the structure of the single crystal of germanium which he holds. 

New frontiers in metallurgy. During the first fifty years of the 
twentieth century, the number of metallic elements used commer- 
cially has more than tripled. Research has heen responsible for this 
great surge in the use of metals. For example, our nuclear energy 
program has resulted in the discovery of new metals such as plu- 
tonium and in the commercial production of such metals as zir- 
conium and beryllium; jet engine research has produced a variety of 
high-temperature alloys using the "new" metals as well as substances 
known as cermets, a combination of metal and glass or other ceramics. 

But despite the advances which have been made, metallurgy re- 
mains a great challenge to man's ingenuity. Fully a third of the 
known metals are waiting for the chemist who can contrive methods 
of producing them in large quantities for the service of mankind. 

Nor is the problem provided by metals whose metallurgy is still 
unsolved the only mystery left. The metals in general use today are 
never pure. Even the so-called commercially pure metals contain 
small but important amounts of impurities. "Pure" gold, for exam- 
ple, contains as much as 0.4 percent copper and other substances. 
Only recently have science and industry begun to explore the prepa- 
ration of truly pure metals in which the impurities may, in some 
cases, be less than one part in 100 million. 



Even a trace of impurity greatly affects the behavior of a metal. 
Tiny crystals of virtually pure iron produced in the laboratory have 
been found to be remarkably flexible and to possess a tensile strength 
of almost a million pounds per square inch, far beyond that of any 
commercial alloy or metal. 

As yet, no simple inexpensive commercial method of producing 
such pure metals has been developed although many are now un- 
der study. Titanium and zirconium, which show the characteristics 
of commercially valuable metals only when they are extremely pure, 
are produced by the Kroll process, named for Wilhelm Kroll who 
developed it for the U.S. Bureau of Mines. The key step in the proc- 
ess is the reduction of a chloride of the metal by molten magnesium 
in an inert atmosphere or vacuum. This process is still not regarded 
as entirely satisfactory. If successful methods of producing pure met- 
als in quantity are discovered, it will undoubtedly be possible for 
metallurgists to create an entirely new series of metal alloys with 
strength, heat and corrosion resistance, and other properties much 
superior to anything now produced. The field of the pure metals 
represents another great frontier in metallurgy to be explored by the 
chemists of the future and offers a worthwhile inspiration to today's 
students of chemistry. 

Metals are unequally distributed over the earth. Metals were dis- 
tributed over the earth without regard for nations. Many countries 
lack essential metals and must import them from foreign sources. The 
United States, for example, must import part or all its supplies of 
antimony, chromium, cobalt, manganese, mercury, nickel, platinum, 
tin, wolfram, and vanadium (see Table 13 below). 

This unequal geographical distribution of the metals has in times 
of peace brought nations closer together in friendly trade. It has also, 
unfortunately, caused some nations to embark upon imperialistic 
ventures, during which regions rich in various natural resources have 
been seized and exploited. It has encouraged conservation and also 
spurred scientists to find substitutes such as plastics for these essen- 
tial metals. Shortages of metals have also led to thorough research in 
the extraction of metals from ores previously considered too poor in 
metal content. 


(In Approximate Percentage of Consumption) 

Mercury 100 Chromite 90 Manganese 82 Wolfram 80 Antimony 44 
Nickel 92 Tin 82 Platinum 81 Cobalt 78 Vanadium 40 


A itody of thi$ table will help you to organize and remember what you have learned about these metals. |fcfc *** *t > * <$K* 

SSfc^V '^ .<&& 

*?***<- ^ 















Mexico, Bolivia, China, U. S. 

U.S., Canada 

Germany, U. S. 

U. S., Mexico 

Turkey, Rhodesia, U. S., Philippine Is. 

Belgian Congo, Canada 

South Africa, U. S. 

Mexico, U. S. 

Available all over the world 

Russia, South Africa, India, U. S. 

Spain, Italy, U.S. 

Canada, New Caledonia, Russia 

Russia, Canada, U. S. 

Mexico, U.S. 

Malay States, Indonesia, Bolivia 

U. S., India, Norway 

China, Burma, U. S., Spain, Portugal 
Peru, U.S. 
U.S., Belgium 


Sb 2 S 3 (stibnite) 

As 2 S 3 (orpiment) 

Free and combined 

CdS (greenockite) 

FeO Cr.O : , (chromite) 

CoAs, (smaltite) 

Free condition 

PbS (galena) 

Sea water, KMgCl a -6H:,O (carnallitej 

MgCO 3 -CaCO 3 (dolomite) 

MnOo (pyrolusite) 

HgS (cinnabar) 

Complex ores 

Free condition 

Free and Ag 2 S 

SnO a (cassiterite) 

FeTiO 9 (ilmenite),TiO 2 (rutile) 

Wolframite and scheelite (CaWOJ 
VjS 3 (patronite) 
ZnS (zinc blende) 

^W^l\X x ?,V" -, 

"^X^i^I^ J^^> 

* ^5?)^, ? < ..X * 

: i%H>'^ 

^M^-:i* ^^ 

^^:j\^ 5 ^ 


*. J <2^ >\* .is 


Roast, reduce 
Roast, reduce 
(mp 271 C) 
Roast, reduce 
Reduction with Al 
Placer, cyanide 
Roast, reduce 


Roast, distill 

Heat with CO, decompose Ni(CO) 4 

(mp 1750 C.) 

Parke's process 

Reduction (mp 232 C.) 

Kroll process 

(mp330(T C.) 
Roast, reduce 


Type metal 


Low mp alloys 

Electroplating, bearings 


High temperature alloys 

Jewelry, coinage 

Batteries, tetraethyl lead 

Structural products 

Steel and aluminum alloys 


Plating, coinage 


Coinage, cutlery 


Pigments, alloys 

Electric bulbs, tool steels 
Tool steels 

Babbitt metal 
Lead shot 
Wood's metal 
Wood's metal 
Chrome steel 

24 carat (pure) 
Type metal, solder 
Aluminum alloys 

Dowmetal G 


Monel metal 



Solder, pewter 

manganese titanium 

chromium titanium 

Tungsten steel 


Bronze, brass 


Chemistry, Jan., 1945, pp. 21-22. Contains a very interesting 
account of several "rare" elements which are plentiful. 

U.S. Department of Agriculture. Painting on the Farm. 
Bull. No. 1452. Supt. of Documents, Washington, D.C. 5f 

Weeks, Mary E. The Discovery of the Elements. American 
Chemical Society, Easton, Pa., 1945. This scholarly and well- 
illustrated book tells the story of the discovery of the elements 
and something about the discoverers. 


Group A 

1. Copy and complete. Do not write in this book. Eight of 
the last elements discovered were produced by Americans. 

The symbols for elements 93-98 are Of all the elements 

. . . are metals. The most important metals used for coating 
other metals are . . ., . . ., ...,,..., . . ., and Au. The chief 
ore of Sn is .... The most common ore of Pb is .... Hg is 
usually obtained from cinnabar, whose formula is .... The 
symbol for wolfram is 

2. What is the difference between rusting and corrosion? 

3. (a) What are the three general methods for protecting 
metals from rusting and corrosion? (b) Why should steel arti- 
cles be oiled thoroughly before being stored? 

4. Compare the process of tin-plating, nickel-plating, and 

5. (a) Tin cans are not made of pure Sn. Why? (b) Name 
two highly magnetic alloys. 

6. (a) What is the composition of our silver coins, (b) of 
gold coins formerly used in the United States, and (c) of ster- 
ling silver? e 

7. What are two reasons for the rapid development and 
widespread use of chromium-plating? 

8. Using an equation, explain the metallurgy of Cr. 

9. Using equations, describe the metallurgy of Zn, starting 
with its sulfide ore. 

10. (a) What metal does Cd resemble? (b) For what is it 
used? (c) Why? 

11. Compare the methods used in the purification of Cu, 
Zn, and Hg. 




12. (a) What are five metals that were known to the 
ancients? (b) Why were these metals known to them? 

13. Explain chemically the formation of tarnish on silver- 

14. Using equations, discuss a chemical method for cleaning 

15. Using a diagram, show how you would silver-plate a 

16. What is placer mining? 

17. How does a jeweler tell pure Au from brass? 

18. What is 14-carat gold? 

19. What are (a) liquid gold, (b) gold leaf, and (c) white 

20. What two properties of Pt make it very valuable to in- 

21. For what two reasons have chemists made an effort to 
find substitutes for Pt? 

22. (a) What is the melting point of W? (b) Name its chief 

23. Compare the properties of the two chlorides of Hg. 

24. (a) How is Hg used in fluorescent lighting? (b) What 
are the advantages of fluorescent lighting? 

25. (a) Describe a recently introduced piece of apparatus 
that makes use of both Hg and W. (b) What are phosphors? 

26. Using word-equations, describe the manufacture of 
white lead. 

27. What types of substances do all paints contain? 

28. (a) Name four elements that appear in the nitrogen 
group in the periodic table, (b) What property have they 
in common? 

29. What effect would the sudden exhaustion of sources of 
Sb have upon one particular industry? 

30. What are three metals in addition to As whose com- 
pounds are poisonous? 

3 1 . How is Pb used in the chamber process of manufacturing 
H 2 S0 4 ? 

32. Using a diagram, describe the construction of a lead 
storage battery. 


33. How does the use of the word carat in connection with 
Au differ from its use in connection with diamonds? 

34. Cu, Ag, and Au appear in the same group of the periodic 
table. Why? 

35. Why are lead paints not used in chemical laboratories? 

Group B 

36. How would you prepare (a) ZnSO 4 , (b) ZnCO 3 , 
(c) AuCl 3 , and (d) As 2 O 3 ? 

37. Why is it dangerous to smoke around batteries that are 
being charged? 

38. The bottoms of ships are often painted with HgO. 

39. What is the function of a flux in soldering? 

40. Pure Zn does not dissolve in dilute H 2 SO 4 . Commercial 
Zn reacts readily with this acid. Explain. 

41. The atomic number of Hg is 80, and that of Au is 79. 
The transmutation of Hg into Au has been accomplished. 

42. For what are the following metals used? (a) Ti, (b) Zr, 
(c) Be, (d) Ge, (e) Co. 

43. (a) Discuss powder metallurgy, (b) For what are the 
products of this process used? 

44. In the fight against metal corrosion the process of 
cathodic protection has lately been introduced. Explain this 


1. Make a collection of ores of as many metals as you can 
find in the community in which you live. What do you con- 
elude from a study of your collection? 

2. Consult your teacher of history or economics and write 
a two- or three-page report on the present gold and silver 
policies of the United States. 

3. Write a report on the way in which you or your father 
or some friend takes care of the lead storage battery in your 
car. Include in your report such items as charging and dis- 
charging, checking, and testing the specific gravity of the elec- 



A farmer with his potash locked 
up in silicates is like the merchant 
who has left the key of his safe at 
home in his oiher trousers. He may 
be solvent, hut he cannot meet a 
single draft. It is only solvent potash 
which plants can use. E. E. Slosson, 

Liebig used chemistry in agriculture. Fertilizers, as you know, 
are composed of substances that must be replenished in the soil to 
prevent plant starvation. Soils have been fertilized from time im- 
memorial with manure, fish scrap, dried blood, wastes from slaugh- 
terhouses, bird excrement (guano) , and bones. The practice of the 
ancients was based not upon chemical knowledge but upon their 
general observation that crops thrived better if grown on land sup- 
plied with refuse or decayed organic matter of any kind. 

Justus von Liebig, who made many contributions to organic chem- 
istry, and whose name is commonly linked with the Liebig condenser, 
is generally regarded as the father of agricultural chemistry. In 1840 
he tested his new theory of soil fertility on a barren piece of land near 
Giessen, Germany. Year after year he kept on feeding the soil with 
only mineral fertilizers, including nitrates, phosphates, and potas- 
sium salts, until it was as fertile a spot as could be found in all Ger- 

With one blow Liebig overturned the firmly rooted belief that 
plants thrive only on manure or other organic matter. He proved 
that they would also thrive on soluble salts containing potassium, 
phosphorus, and nitrogen. He showed that ' 'since plants assimilate 



potassium, nitrogen, and phosphorus, these elements must neces- 
sarily be resupplied to the soil/' Other men followed this pioneer 
work. One of these investigators, Sir John Lawes, started the Rotham- 
sted Experiment Station at Harpenden, England, now the most fa- 
mous agricultural experiment station in the world. 

The fight to save topsoil. Erosion and an exhaustive system of crop- 
ping have taken terrific toll of the soil in many parts of the United 
States. Fifty million acres of land in this country have been so ruined. 
In part, this has been caused by the destruction of forests and the 
poverty of some farmers who were forced to "eat up" their soil by 
not returning to it the elements necessary for its fertility. 

The practice of contour farming, that is, farming around a hill or 
slope and not in straight lines, holds water as it falls, and helps to 
prevent further erosion. Strips of close-growing crops, such as grasses, 
planted between other strips of cultivated crops, such as wheat and 
corn, also help to hold the topsoil. This is known as strip cropping. 
Plowing under green-manure crops, such as alfalfa and sweet clover, 
helps both to keep rich soil in good condition and to improve poor 
soil. In addition, the generous use of irrigation and commercial fer- 
tilizers is essential to the success of farming. 

Krilium is the trade name for a number of soil conditioners which 
help erosion control by making clay-like soil porous and friable. It 
also raises its water-holding capacity and improves root penetration. 
It is a synthetic, negatively-charged particle which is adsorbed on 
clay forming small lumps which bind the soil together, and cannot 
be washed out like most natural soil conditioners. 

Potassium salts for fertilizers. Plants require sulfur, iron, calcium, 
magnesium, manganese, silicon, copper, boron, zinc, and minute 

Fig. 1O2. Topsoil lost through 
erosion. Topsoil carried away 
by erosion cannot be recovered. 
How may erosion be retarded? 



quantities (traces) of other elements for growth. However, the three 
elements most often lacking in sufficient quantities are nitrogen, 
phosphorus, and potassium. For example, if potassium salts in the 
soil in which potatoes have been planted are insufficient, the result- 
ing crop will illustrate very strikingly the potash hunger of the po- 

Our potash problem. In early days, because our land was fertile, 
there was no potash problem. In fact, for many years the principle 
export commodity of the American Colonies and the newly formed 
United States was crude potassium carbonate, K 2 CO :i , obtained by 
washing out (leaching) this chemical from wood ashes in pots by 
means of water. This industry started as early as 1608 in Jamestown, 
Virginia, and continued until rich European deposits were discov- 
ered about the time of the War Between the States. At this time huge 
deposits of potassium salts began to be developed from the salt mines 
at Stassfurt, Germany. Here in layers hundreds of feet thick, is an 
almost inexhaustible supply of soluble potassium in the form of 
sylvite, KC1, and carnallite, KC1 MgCL 

When World War I brought home to us our extreme dependence 
upon foreign countries for our supplies of potassium salts, we began 
to look for domestic sources, and found them. In the ^5 years be- 
tween World Wars I and II, the combined efforts of the U.S. Gov- 
ernment and private agencies succeeded in building up an industry 
that freed us from foreign supplies. 

Our main sources are the huge underground deposits of KC1 and 
K 2 SO 4 in Carlsbad, New Mexico, and Texas. Searles Lake, an inland 
sea in the Mohave Desert, contains 36 percent salt of which one-fifth 
is KC1. Other minor sources are the ashes of kelp (a giant seaweed 

This demonstration plot shows how fertilizers help crops grow and halt erosion. 

U,8. Department of Agriculture 

' ' ' 

A stockpile of ammonium 
sulfate which can be used 
either as a fertilizer or in 
the manufacture of other 
ammonium salts. 

off the coast of California which also contains iodine) , sugar beet 
wastes, basic lining of steel furnaces, and the dust of cement works. 

The future of our fertilizer industry. The fertilizer industry is one 
of our basic industries. Today the fertilizer manufacturer, dependent 
more and more upon mineral sources, is learning to compound fer- 
tilizers from them in order to meet the requirements of each kind of 
crop. To save freight charges, these fertilizers are being put on the 
market in a highly concentrated condition. 

"Ammophos," ammonium phosphate, made by combining ammo- 
nia, NH.p and phosphorus pentoxide, P 2 O., is such a fertilizer. Some 
others, such as potassium ammonium phosphate, contain all three of 
the elements necessary in a complete fertilizer. A complete fertilizer 
for sugar beets now on the market is a 4-12-6 fertilizer. This means 
four percent nitrogen, 12 percent PX),,, and six percent K 2 O. 

Will feldspar be a future source of potassium? One of the most 
common rock-forming minerals in this country is feldspar, composed 
of potassium aluminum silicate, KAlSi :t O s . This mineral is a tremen- 
dous untapped storehouse of potassium. Partly because feldspar is 
insoluble in water, chemists have not yet succeeded in getting its 
potassium into soluble form inexpensively. A solution to the problem 
would give us an inexhaustible supply of potassium, and also bring 


> . 'M^'.'V , >*% ,.y 

| Nitrogen E3 Potassium 3 Magnesium 
| Phosphorus Calcium EO Sulfur 




Adapt*! from (1. S. 0porfmnt of 




Ai'mftontout Publkvthn No. 400 

a fortune to the man who finds it, Incidentally, feldspar is rich also 
in aluminum, but thus far no one has yet been able to devise a 
method to extract the metal from it at a cost low enough to be com- 
petitive with the Hall process. 

Potassium nitrate, KNO S , chief constituent of gunpowder. With 
the introduction of gunpowder, potassium nitrate became a much 
needed chemical. Gunpowder is a mixture of carbon, sulfur, and 
potassium nitrate. When ignited, tremendous volumes of gases and 
vapors are produced suddenly, thus giving it explosive power. The 
high temperature at which the reaction occurs increases these tre- 
mendous volumes enormously. The equation for the explosion of 
gunpowder may be written: 

2KNO 3 + 3C + S - K 2 S 4- N 2 + 3CO 2 

Sodium nitrate cannot be used instead of potassium nitrate in gun- 
powder because it is hygroscopic, that is, it takes up water from the 

Potassium nitrate is found jn the topsoil of certain sections of 

tdia, Persia, and other countries, but the supply is not as great as 
at of sodium nitrate. From the timenhat gunpowder began to be 
used in Europe, the constant warring of nations led to a great de- 
mand for this chemical. Among the first papers Communicated to the 
English Royal Society was one that dealt with the manufacture of 



saltpeter. Its author declared "the only place where saltpeter is to 
be found is in stables and pigeon houses/' 

The East India Company imported nitrate from India as early 
as 1625. In France and Germany "niter plantations" sprang up. 
These were heaps of organic material mixed with wood ashes and 
other alkalies. The putrefaction that went on in these heaps in the 
open air produced potassium nitrate as one of the products. Up to 
the time of Lavoisier, officers of the French government had been 
accustomed to search for niter in wastes in the cellars of private 
houses, and to confiscate it, both to help supply the needs of the 
country and to prevent its use in insurrections. But shortly before the 
French Revolution, Lavoisier, as controller of munitions, abolished 
this confiscation and at the same time increased the supply of this 
chemical by improving methods of manufacture. 

About 1650, Glauber showed that saltpeter could replace manure 
as a means of restoring the fertility of exhausted soils. It is still widely 
used as a fertilizer. Today it is prepared from sodium nitrate by the 
addition of potassium chloride solution, as follows: 

" ' NaNO 3 + KC1 ?=i NaCl + KNO 3 

Although this reaction is reversible, the potassium nitrate can be 
separated out by fractional crystallization (see page 473) because 
potassium nitrate is more than six times as soluble as NaCl in boil- 
ing water. 

Will hydroponics supply our plant foods in the future? Hydro- 
ponics (from hydro, water, and ponos, labor) is a method of growing 
plants in water to which chemicals are added rather than in soil. 
This growing of crops without soil with the aid of the proper plant 

Apparatus for fertilizing the soil by injection of anhydrous ammonia. 

Shell Chemical Corporation 


nutrients will in some cases increase the yield and improve the prod- 
uct. Tomatoes, beans, cucumbers, and many other plants have already 
been grown successfully by hydroponics. 

It is not likely that science will do away completely with farms, 
such as we know, in favor of smaller areas filled with tanks and sterile 
sand beds containing the solutions required to produce various crops. 
However, much about plants can be learned by using hydroponics 
as an experimental tool of the agrobiologist. 

What properties are common to the salts of the alkali metals? 
Chemists have found that certain characteristics of the salts of the 
members of a family of elements resemble one another. Hence, they 
study the salts of such a family together. 

We have already learned about some of the salts of potassium and 
shall continue with the salts of sodium, another important member 
of the alkali metals. The alkali metals form part of Group I of the 
periodic table. All their salts are white, crystalline solids. They are all 
soluble in water. Potassium compounds are just about as abundant 
as sodium compounds. 

After studying the salts of sodium and potassium you will notice 
their marked resemblance to the ammonium (NH 4 ) salts. Because 
of this resemblance and also because NH 4 + is a univalent, positively 
charged ion, the NH 4 ion might be considered as belonging to the 
same family group as sodium and potassium the alkali metals. 
Because ammonium hydroxide is a base, and at the same time read- 
ily gives off a gas, or volatilizes, it is often called volatile alkali. 

Sodium nitrate, fertilizer from the Andes. The coastal mountains 
in northern Chile rise abruptly from the floor of the Pacific Ocean. 
Between this range and the westward slopes of the Andes Mountains 
at elevations ranging from 4000 to 9000 feet is a desert plateau where 
only occasional cloudbursts send water into the loose, sandy soil. 
Most of this plateau is as rainless as any part of the world. In certain 
sections, annual rainfall is less than an inch. 

Below this sandy soil is caliche, a rock containing between 50 per- 
cent and 75 percent sodium nitrate, NaNO 3 . This nitrogen-contain- 
ing mineral probably was formed by the action of nitrifying bacteria 
on ancient accumulations of seaweeds, guano, and other organic ma- 
terials as well as by ele^|Pi<^fccation during electric storms. The de- 
posits were preserved b*^$|| o the scarcity of rain. If the rainfall 
were normal, the deposits would hav^v dissolved and washed into 
the Pacific Ocean. Another theory explains that these deposits have 
come from volcanic ash rich in nitrates or accumulated by the minute 
nitrate content of the underground \Vatrs of the region. 

Workmen in the Chile ni- 
trate deposits using pneu- 
matic drills to break up 
large chunks of caliche. 

Chilean Nitrate Educational Bureau, Inc. 

Here in the Andes is the greatest concentration of nitrate salts in 
the world, a circumstance which for many years gave Chile a world 
monopoly of nitrate fertilizer. This monopoly was broken by the 
development of the synthetic nitrate industry (see page 262) . Chile 
still continues to export this essential chemical in large quantities. 

We are no longer dependent upon Chile for our nitrates, and in 
normal times we actually export some synthetic sodium nitrate to 
manufacturers and farmers both in Europe and Asia. Nitric acid is 
first synthesized from the elements of the air, and then the HNO 3 is 
treated with soda ash, forming synthetic NaNO 3 . 

2HNO 3 + Na 2 CO 3 - 2NaNO 3 + H 2 O + CO 2 

Chile saltpeter, another name for sodium nitrate, is used chiefly as 
a nitrogen fertilizer, and also in making nitric acid and potassium 
nitrate. A most important byproduct of caliche is iodine, extracted 
from the sodium iodate, NalO.,, found mixed with sodium nitrate. 

Sodium chloride, the best-known salt. Whenever most people 
speak of salt, they mean sodium chloride, NaCl. Not only has so- 
dium chloride been known from antiquity, but even today it is one 
of the most abundant and important chemicals in the service of man. 

The sea is an inexhaustible salt supply. If the salt of all the oceans 
could be collected, it would occupy 15 times the volume of Europe 
above sea level. Sea water contains an average of about three percent 
solid matter, most of it sodium chloride. The Dead Sea contains 




seven percent sodium chloride and 13 percent of another salt, 
magnesium chloride. The presence of this latter salt (because of its 
deliquescence) causes table salt to cake in moist weather. Prepared 
salt often contains 0.1 percent of tricalcium phosphate or one per- 
cent CaCO 3 to keep it free-flowing. 

It is too costly to separate salt from sea water by boiling. Therefore 
sea water is usually collected in shallow basins near the shore, and 
the salt is concentrated by evaporation resulting from the heat of the 
sun, or solar evaporation. Early colonists in America obtained their 
salt by this method. During the winter in some countries, such as 
China and Japan, sea water is run into shallow troughs, and as the 
top layers of water freeze, the ice is removed until the salty sea wa- 
ter below becomes very concentrated. This concentrated brine is then 
heated to drive off the remaining water. 

Salt is obtained also from salt lakes, as in India and at the Great 
Salt Lake in western United States, and from deposits of solid salt 
known as rock salt, or halite. Millions of years ago the sea covered 
part of what is now dry land. As the land rose, or the sea receded, 
or both, sea water was trapped in depressions in the land. This wa- 
ter gradually evaporated, leaving deposits of solid layers of salt. This 
explanation of the origin of the rock-salt mines near Syracuse, New 
York, and the famous potassium salt deposits of Carlsbad, New Mex- 
ico, is accepted generally. 

Rock salt, containing as much as 95 percent sodium chloride, is 
either mined like coal or brought to the surface by first dissolving it 

Dow Chemical Company 

Brine flowing from salt 
wells into a storage tank. 
Great quantities of salt are 
essential in industry par- 
ticularly in the manufacture 
of chemicals. 


in water forced down through pipes, and then pumping the satu- 
rated salt solution to the surface by compressed air. Most of the salt 
used in this country is obtained from brine wells. 

Though impure salt has many uses, for certain other uses it must 
be purified. To accomplish this, hydrogen chloride gas is bubbled 
through a concentrated solution of impure salt, causing pure sodium 
chloride to crystallize out (see common-ion effect, pages 244-245) . 

Fractional crystallization separates a mixture of salts. The various 
salts found mixed with sodium chloride have different solubilities in 
water. Their separation from one another depends upon both their 
solubility in water and their concentrations. A less soluble salt 
crystallizes out before a more soluble salt, provided their concentra- 
tions are the same. A salt present in a higher concentration crystal- 
lizes out before a salt present in a lower concentration, provided their 
solubilities are the same. Hence, that salt which is present in solu- 
tion in a relatively concentrated form and whose solubility is com- 
paratively low is the first to separate from a solution containing a 
number of different salts. This is the principle of fractional crystal- 
lization, used frequently in the preparation of pure salt crystals such 
as sodium bicarbonate and potassium nitrate. 

How sodium chloride is used. Salt plays an important part in the 
bodies of all animals, because it is necessary in maintaining the wa- 
ter content of the tissues. Plants and animals originally developed in 
sea water. The blood of most animals resembles sea water. The ani- 
mal maintains its cells in an internal environment resembling the 
ancient sea of its origin. Blood has a fixed amount of salt; to reduce 
this amount very much results in throwing the normal processes of 
the body out of balance. In fact, in cases of shock following injuries 
or surgical operations, salt in the form of normal saline solution, 
0.8 percent salt in water, the concentration of salt in blood serum, 
is injected into a vein or under the skin. 

So necessary is salt to animal life that wild herbivorous animals 
are known to travel great distances in order to reach "salt licks" 
where they can satisfy their craving for this mineral. Human beings 
use salt as a seasoning in almost all foods. This salt, with the salt 
present naturally in foods, supplies the teaspoonful of salt that an 
adult needs daily. From it, gastric cells produce the hydrochloric acid 
present in the gastric juice. In conditions of excessive perspiration, 
great quantities of salt are lost from the body, and should be re- 
placed by a greater-than-normal salt intake. Salt, in the form of salt 
tablets or salt from other sources, may be used to prevent the effects 
of this kind of salt starvation. 


In addition to its use as a food, salt is a key chemical for it serves 
as the foundation material for numerous industrial processes. Of 15 
million tons of salt consumed annually in this country, only about 
one pound out of ten is used in foods. The remaining nine pounds 
are used in many industrial processes such as the manufacture of 
NaOH, NaHCOg, HC1, Na, C1 2 , soap, glass, and enamel. 

Rock salt is used in curing fish, meat packing, curing hides, and 
in making freezing mixtures, especially for the manufacture of ice 
cream. This use is based on the principle that a salt, as it goes into 
solution, absorbs heat from any material with which it is in contact. 
In addition, melting ice absorbs heat. Salt mixed with ice or with 
water near its freezing point will produce a temperature as low as 
22C. For years, salt has been used in controlling ice on city streets 
and on the subsurface soil of the beds of railroad tracks to prevent 
the tracks from heaving as a result of freezing. Sodium chloride is 
used also in water-softening (see page 491) . 

Salt on the farm. Great quantities of salt are consumed on farms. 
Rock salt and salt blocks are used to provide artificial salt licks for 
farm animals, and salt is added to commercially produced animal 
feeds. Farm food preparation, including the canning and preserving 
of vegetables, fruits, and meats, requires salt. Weed killers and insec- 
ticides often contain salt, and many commercial fertilizers require 
salt in their production. 

The economic importance of salt is very great. Some nations keep 
the production and sale of salt a government monopoly or at least 
levy a tax on it. In times past, cakes of salt have even been used as 
money, as, for example, in Tibet and Abyssinia. In our own coun- 
try, during the War Between the States, one of the main purposes 
of a Union campaign into southwestern Virginia was to capture and 
destroy one of the chief sources of salt of the Confederate States, 
located at Saltville. 

The manufacture of sodium bicarbonate, chief constituent of 
baking powder. Sodium hydrogen carbonate, or sodium bicarbonate, 
NaHCOg, is manufactured chiefly by the Solvay process, which re- 
placed the older Le Blanc process invented in France about the 
time of the French Revolution. Ernest, at 23, and Alfred Solvay, 
Belgian chemists, perfected the new process about 1863. In this proc- 
ess purified brine is sprayed in at the top of an absorption tower, in 
which it meets ammonia gas and carbon dioxide. The following re- 
action takes place. 

NaCl + H 2 O + NH 3 + CO 2 - NH 4 C1 + NaHCO 3 


The temperature must be carefully controlled, since otherwise this 
reaction is reversible. The sodium bicarbonate produced is separated 
from the ammonium chloride by fractional crystallization. 

The ammonium chloride is heated with slaked lime to recover the 
ammonia gas, which is used again. 

Ca(OH) 2 + 2NH 4 C1 -> 2NH 3 1 + CaCl 2 + 2H 2 O 

The carbon dioxide used in the Solvay process is obtained from 
limestone. One of the largest plants in which sodium bicarbonate is 
manufactured by this process is located at Syracuse, New York, where 
the raw materials are nearby and plentiful. 

Sodium bicarbonate is a constituent of all baking powders and is 
therefore called baking soda. Under the name bicarbonate of soda it 
is used in medicine. A very effective, yet low-cost, dentifrice can be 
prepared by mixing one part by weight of sodium bicarbonate with 
approximately two parts of salt. 

Sodium carbonate, or washing soda. Most sodium bicarbonate is 
converted into sodium carbonate, Na 2 CO 3 . When sodium bicarbonate 
is heated, carbon dioxide is liberated and anhydrous sodium car- 
bonate, a white powder, is formed. The carbon dioxide formed in 
the reaction is used in the Solvay process. 

2NaHCO 8 -> Na 2 CO 3 + H 2 O + CO 2 1 

When this sodium carbonate is dissolved in water and crystallized 
out, the crystals formed contain ten molecules of water of crystalliza- 
tion, and are called sal soda. 

Na 2 CO 3 + 10H 2 O Na 2 CO 3 - 10H 2 O (crystallized washing soda) 

This salt, Na 2 CO 3 10H 2 O, is efflorescent and, on exposure to air, 
loses nine molecules of water, forming the stable white salt, 
Na 2 CO 3 . H 2 O. Upon heating, this stable salt yields Na 2 CO 3 . Soda 
ash obtained from burning seaweed contains about 98 percent so- 
dium carbonate. 

Because sodium carbonate is the salt of a strong base and a weak 
acid, it hydrolyzes in part, forming sodium hydroxide (lye) . For this 
reason, sodium carbonate is a good cleansing agent. And because 
sodium carbonate is a valuable water-softener (see page 489) , it is 
used to some extent in home laundering and in laundries as washing 
soda. The recently-developed "Ammo" is a mixture of Na 2 CO 8 and 
an ammonium salt. 

Soda ash is industry's third most important chemical. More than 
4.5 million tons of it are produced annually in the United States. 


The manufacture of glass consumes one-third of this amount, and 
an equal amount goes into the manufacture of soap and other cleans- 
ers. It is used also in tanning, as a lumber dip to prevent formation 
of stains that are produced by fungi and molds, and in the manufac- 
ture of paper pulp. 


Laurie, Alexander. Soilless Culture Simplified. McGraw-Hill 
Book Co., New York, 1940. 

Mansfield, George R. "American Potash Reserves." Indus- 
trial and Engineering Chemistry, Dec., 1942, pp. 1417-1421. 
Deals with the potash industry in the United States and dis- 
cusses also Federal and private development of our resources of 
potassium compounds. 

Moore, F. J. History of Chemistry (3rd ed.) . McGraw-Hill 
Book Co., New York, 1939. The life and work of Liebig and his 
friendship with Woehler. 


Group A 

1. The ancients fertilized their fields, (a) What method 
did they use? (b) Is it likely that they understood the chem- 
istry of this method? (c) Why? 

2. Copy and complete the following statements. Do not 
write in this book. Chemistry was first used in agriculture 
by .... He showed that plants could live upon . . . independent 
of organic matter such as manure. In England a man named 
. . . started the first agricultural experiment station at .... 

3. (a) Why should most soils be fertilized? (b) What three 
elements must be in fertilizers? (c) Why? 

4. What are three methods of preserving the topsoil? 

5. How was K 2 CO 3 obtained by American colonists? 

6. Is America still dependent for its potassium compounds 
upon foreign sources? Explain. 

7. What is a complete fertilizer? 

8. (a) What mineral plentifully distributed in the earth's 
crust is a compound of K? (b) What problem must be solved 


before this mineral can be used as a source of potassium 

9. Name four sources of potassium salts in the United 

10. What are the chief sources of (a) NaNO 3 and (b) 
KN0 3 ? 

11. Compare the action of KNO 3 on litmus with that of 
Na 2 CO 3 . Explain. 

12. (a) What element used often as an antiseptic is found 
with NaNO 3 in caliche? (b) In what compound does it occur? 

13. What were niter plantations? 

14. (a) What is the composition of gunpowder? (b) Why 
cannot NaNO 3 be used instead of KNO^ in gunpowder? (c) On 
what does the explosive quality of gunpowder depend? 
(d) Write the reaction for the explosion. 

15. A manufacturer has 200 tons of KNO 3 on hand. What 
is the greatest weight of gunpowder that he can prepare? 

16. Which is the better fertilizer: NaNO 3 or KNO 3 ? Ex- 

17. Write a balanced equation showing how KNO 3 can be 
made from NaNO 3 . 

18. (a) What is hydroponics? (b) How widely used may 
hydroponics become? (c) Why? 

19. In what three ways do most sodium, potassium, lithium, 
and ammonium salts resemble one another? 

20. (a) What are the chief sources of NaCl? (b) Why does 
common table salt become lumpy? 

21. (a) NaCl is a necessary constituent of the food of all ani- 
mals. Why? (b) In certain countries all NaCl is taxed. Why? 

22. (a) What is fractional crystallization? (b) Illustrate it. 

23. Describe two methods used in extracting NaCl from sea 

24. (a) What is rock salt, and (b) how was it probably 

25. (a) What four important chemicals use large quanti- 
ties of NaCl in their manufacture? (b) What are six other 
uses of NaCl? 


26. Write equations illustrating the preparation of NaCl 
by: (a) direct combination, (b) neutralization, (c) the com- 
bination of a carbonate and the proper acid, (d) The last two 
reactions go to completion. Why? 

27. Make a table of three columns. In the first column 
write: (a) baking soda, (b) washing soda, (c) soda ash, and 
(d) bicarbonate of soda. Write the formula for each of these 
four substances in the second column. In the third column 
write the uses for each of these compounds. 

28. Write a single chemical equation that gives the chief 
reaction involved in the Solvay process. 

29. (a) Name the only substance used in the manufacture 
of Na 2 CO 3 that is not recovered and used again, (b) Explain 
how each of the other substances used is recovered and used 

30. Most of the sodium bicarbonate manufactured is con- 
verted into washing soda, (a) Write an equation to show this 
change, (b) Under what condition does the reaction occur? 
(c) How may it be reversed? 

31. Explain what happens when crystals of washing soda are 
left exposed to the air. 

32. Why is it more economical to use powdered, rather than 
crystalline, washing soda? 

33. What are the two chief uses of Na 2 CO 3 ? 

34. Write an equation showing the reaction between baking 
soda and an acid. 

35. Which will liberate more CO 2 gas when acted upon by 
H 2 SO 4 : 200 g. of anhydrous Na 2 CO 3 or 200 g. of NaHCO 3 ? 

Group B 

36. Why is the Nile River valley so fertile? 

37. How can the elements that are removed from the soil 
by certain crops be determined? 

38. How was Na 2 CO 3 produced before the discovery of 
methods of preparing it from NaCl? 

39. According to the electron theory, explain the structure 
of NaCl. 

40. How can you tell the difference between NaHCO 3 and 
Na 2 CO 3 ? 

41. Why does 1 g. of NaOH neutralize more acid than 1 g. 
of KOH? 


42. Describe how KNO 3 is prepared by fractional crystal- 

43. (a) How was Germany helped during World War I by 
the perfection of the Haber process for making synthetic ni- 
trates? (b) In what condition is the United States with regard 
to nitrates? 


1. Write a two- or three-page report on the way in which 
Justus von Liebig's contributions to agriculture have affected 
our ways of living. 

2. If you live in a farm area, make a study of the nature of 
the soil in that area and the types of fertilizers used. Report 
your findings to the class. What do you conclude from your 

3. In spite of all the knowledge that science has accumu- 
lated concerning the proper growing of crops, why do crops 
not flourish in many cases? List as many reasons as you can 
find, and place them in the order of their importance. 

4. Make a study of all the sodium and potassium salts that 
are used in your home and on your farm (if you live on one) . 
State the uses in each case. What do you conclude concerning 
these salts? 

5. Prepare some pure KNO 3 crystals by dissolving pure 
NaNO 3 and pure KC1 in a minimum volume of hot water (use 
about 10% more NaNO 3 than KC1) . After the solution cools 
down to room temperature place it in a refrigerator for an 
hour. Now pour off the liquid and examine the KNO 3 crystals 
formed. Why did the KNO 3 rather than the NaCl crystallize 
out? (Examine solubility curve on page 219.) 

6. Make an investigation of the relative amounts of fer- 
tilizers used in your community during the past 30 or 40 years. 
What do you conclude? 



. . . The water which drips from the 
stalactite, deposits some more calci- 
um carbonate on the floor and, in 
the course of ages, there rises up to 
meet the descending stalactite a cor- 
responding column from the floor. 
Ultimately the two meet and slowly 
coalesce to form those mighty snow- 
white pillars which hold up and sup- 
port the roofs of the most stupendous 
caverns. Geoffrey Martin, 1911 

Calcium and the other alkaline earth metals. Because calcium 
stands high in the replacement series of the metals, it never occurs 
free. Calcium compounds are very stable. Calcium lends electrons 
easily; combines readily with oxygen, water, and acids; and, like 
sodium, is prepared by the electrolysis of one of its fused salts, cal- 
cium chloride. Since the atomic weight of calcium is 40, the struc- 
ture of its atom is represented as in Fig. 103. 

In addition to calcium, Group II of the periodic table includes 
barium, beryllium, magnesium, strontium, and radium the so- 
called alkaline earth metals. When heated in a flame, a barium or 
a strontium compound gives a color to the flame. The strontium flame 
is red; the barium flame is green. Hence, compounds of barium and 
strontium are used in making tracer bullets, fireworks, and signal 
lights, or fusees. 

Calcium carbonate and its forms. Calcium carbonate, CaCO 3 , 
occurs in four forms: limestone, marble, chalk, and calcite. Calcium 
carbonate is the second most widely distributed mineral on earth. 
It is a part of the great rocky masses of sedimentary land formations. 
Limestone is often found mixed with ores and in their metallurgy 
it must be eliminated as gangue. 




Limestone is a sedimentary, or alluvial, deposit of calcium car- 
bonate built up from the shells of minute aquatic animals. Thou- 
sands of tons of limestone are used each year in the manufacture of 
glass. The metallurgical industries consume tremendous amounts of 
limestone as a flux. It is also the sole source of commercial lime. 

Marble is limestone that has been so heated and compressed in the 
ground by natural forces that its close texture permits a fine polish. 
Limestone and marble are used as building materials. The exteriors 
of the buildings in Rockefeller Center in New York contain 150,000 
tons of Indiana limestone. Some of the best marble is used in sculp- 

Chalk is calcium carbonate that was deposited on the sea floor. In 
general, it consists of the shells of certain minute aquatic organisms. 
Pearls, coral, and many shells also are largely calcium carbonate. 
Eggshells are composed almost entirely of CaCO 3 . Hens gets calcium 
in the form of oystershell or grit in their diet. 

Pure crystalline calcium carbonate is colorless and transparent. It 
is called calcite, or Iceland spar. Some calcite crystals measure as 
much as three feet across and are easily split, or cleaved, in three 

How limestone caves are formed. Calcium carbonate is almost in- 
soluble in pure water, but in water containing a good deal of dis- 
solved carbon dioxide, the insoluble calcium carbonate forms soluble 
calcium bicarbonate. 

CaCO 3 + H 2 O + CO 2 > Ca(HCO 3 ) 2 (calcium bicarbonate) 

Carbonic acid in running underground water, therefore, causes the 
gradual dissolving away of limestone deposits, producing caves. 

A stream rich in dissolved calcium bicarbonate often finds its 
way through the roof of a limestone cave. As a drop of this water 
remains suspended from the roof of the cave, it may evaporate, and 

Fig. 103. Structure of the calcium 

The atomic number of calcium is 20 
Calcium is a typical metal 
The valence of calcium is + 2 
Calcium is very active 

Fig. 104. A calcite crystal separates 
light rays so that objects appear 



the calcium bicarbonate may decompose, liberating carbon dioxide 
and leaving behind a crystal of calcium carbonate. 

Ca(HCO 3 ) 2 - CaCO 3 1 + H 2 O + CO 2 1 

The gradual accumulation of millions of these crystals forms an 
icicle-shaped stalactite. 

Another drop may fall to the floor of the cave. After the water 
evaporates, the calcium bicarbonate decomposes, leaving a crystal 
of calcium carbonate on the floor. This may grow into a stalagmite. 
By the joining of many stalactites and stalagmites, columns are formed 
and eventually the cave may be completely filled. The beauties of 
the Luray Caverns of Virginia, the Mammoth Cave of Kentucky, 
the Carlsbad Caverns in New Mexico, and many others owe their 
origin to this chemical action. 

Making lime from limestone. In remote times lime, a white, soft 
solid, was thought to be an element. However, in 1808, Davy showed 
it to be calcium oxide, CaO. It is prepared by heating limestone, 
which liberates carbon dioxide and crumbles to a powder, CaO. 

CaCO, -> CaO 4- CO 2 1 

This reaction is sometimes used as a source of carbon dioxide, in 
which case the lime produced is a byproduct. 

The time-honored way of making lime was to dig a large pit in 
the side of a hill making a draft-hole at the bottom. The pit was then 
filled with material containing calcium carbonate. Fuel was added, 
lighted, and left to burn out. 

National Park ft< i 

This beautiful limestone formation in Carls- 
bad Caverns, New Mexico, is known as "The 
Temple of the Sun." 

Steel lined 
with firebrick 

Fig. 105. Simplified cross section of a 
vertical lime kiln. 

Today the industrial manufacture of lime is carried out in large 
vertical or horizontal rotary kilns. Through the upper end crushed 
limestone from the quarries is added, and through the other end hot 
gases are forced. The temperature in a lime kiln is about 1000C. 
The heat decomposes the limestone, and the carbon dioxide gas is 
forced out of the kiln. 

In the limelight. Lime, also called quicklime., does not burn when 
heated even to a high temperature, but it glows with a brilliant white 
light. Hence limelights were used at one time for stage spotlights and 
footlights. From this use of lime comes the expression "to be in the 

Because lime has an extremely high melting point, it is used as a 
refractory. As you already have learned, lime is used as the basic lin- 
ing in various metallurgical furnaces such as the open-hearth. It is 
used also in the manufacture of iron, because it combines so readily 
with impurities, forming a slag which can be removed. It is used for 
a similar purpose in other industries. 

CaO + SiO 2 - CaSiO 3 (calcium silicate slag) 

What is slaked lime? You may have watched workmen adding 
cold water to white lumps of lime until the mixture seemed to come 
alive (hence, the name quick lime) ; it began to swell, crumble, boil, 
and give off large volumes of steam. Probably you were amazed at 



the sight of cold water raising the temperature of the lime to the boil- 
ing point of water, instead of cooling off the lime, as might be ex- 
pected. This evolution of heat is caused by the exothermic reaction 
of water and lime, a basic anhydride, forming calcium hydroxide, or 
slaked lime. Its "thirst" for water was slaked. 

CaO + H 2 O -> Ca(OH) 2 + heat 

Slaked lime, or Ca (OH) 2 , is a white solid. It is slightly soluble in 
water, forming limewater. Limewater is a fairly strong base. Even 
though calcium hydroxide is only slightly soluble in water, its solu- 
bility decreases further on rise of temperature. It is extensively used 
wherever a low-cost base is required. For example, slaked lime is used 
in removing hair from hides. When slaked lime is added to water in 
large amounts, a white suspension is formed, which we know as white- 
wash, or sometimes, as milk of lime. 

The fertility of soil depends in part upon calcium ions. Calcium, 
though rarely lacking in most soils, is an essential element in plant 
nutrition. Scientists think that it has some connection with the for- 
mation of the cell wall. The legumes are lime-loving plants and re- 
quire large amounts of it. 

Some soils turn blue litmus pink, and exhibit other acid properties. 
Some crops including pineapples, potatoes, and cranberries grow best 
on such soils; others, including sugar beets, clover, alfalfa, and pea- 
nuts, do not thrive on acid soils. Such soils must be sweetened. Their 
sourness may be removed by the addition of finely ground limestone, 
slaked lime, or crushed sea shells. These substances neutralize the 
acids, and this procedure is called liming the soil. 

Scientific farming involves determining the pH value (see page 
235) of a soil before calculating the amount of calcium compounds 
to be added. Lime also acts mechanically, loosening clay soils and 
making sandy soils more firm. 

Ground limestone and other calcium carbonates may be spread on 
the soil at any 'time; but it is usually not considered advisable to 
apply limestone, manure, and ammonia compounds at the same time. 
The limestone must be thoroughly mixed with the soil. An effective 
method of accomplishing this is by harrowing the soil shortly after 
spreading the limestone. Application of calcium compounds other 
than carbonates requires extreme care. For example, lime will pos- 
sibly injure tender plants and seeds if it comes in contact with them. 

Lime in building. If sand is added to freshly slaked lime, a plastic 
mass called mortar is formed. Mortar adheres to wood, bricks, and 
stone, and is used to hold these materials together. 


Some of the chemical and physical changes that take place when 
mortar is setting are still not thoroughly understood, but we do know 
that, as the mass dries and hardens, water is given off and calcium 
silicate is formed: 

Ca(OH) 2 + Si0 2 - CaSiO 3 . 
Slaked lime 4- sand > calcium silicate 

Some of the grains of sand are held in colloidal suspension (see 
Chapter 38) , giving a definite body and hardness to the mass. The 
outside of the mortar comes in contact with the carbon dioxide in 
air, and minute crystals of calcium carbonate are formed, which help 
in the solidification. 

Ca(OH) 2 + C0 2 - CaCO 3 + H 2 O 

In most construction work, lime mortar has been largely displaced 
by the much stronger cement, but it is still widely used in home 
building (see Chapter 33) . 

If quicklime remains exposed to air for a long enough time, it 
reacts with water vapor in the air, forming calcium hydroxide. By 
uniting with the carbon dioxide in air, this calcium hydroxide forms 
calcium carbonate. Quicklime to which this has happened is said 
to be air-slaked. Such air-slaked lime cannot be safely used for mak- 
ing mortar. Why? 

How slaked lime is used in preparing bleaching powder. Bleaching 
powder, CaOCL, already discussed in the chapter on chlorine, is pre- 
pared commercially by passing chlorine vapor through rooms con- 
taining a three-inch layer of freshly slaked lime. A grayish-yellow 
powder called chloride of lime or chlorinated lime is formed. 

Ca(OH) 2 + C1 2 -> H 2 + CaOCl 2 

A more recent method combines free chlorine and lime directly as 
they mingle with each other while passing through a pipe. 

CaO -I- C1 2 -> CaOCl 2 

Bleaching power is not as pure or stable as Ca(CIO) 2 . It liber- 
ates its chlorine when heated or when acted upon by even weak acids, 
such as the carbonic acid in the air. 

CaOCl 2 + H 2 CO 3 -> CaCO 3 + H 2 O + C1 2 1 

Bleaching powder thus furnishes a simple means of obtaining free 

Spreading plaster of Paris on a form. 
After the plaster hardens, it will be 
removed and used as a mold. 

chlorine for bleaching and disinfecting purposes. It is used also in 
the manufacture of paper. 

Calcium sulfate and plaster of Paris. The great blocks of the 
pyramids of Egypt were cemented together by a plaster made from 
a powder obtained by heating a white, soft rock called gypsum, 
CaSO 4 2H,O. This material, still used today, is called plaster of 

To make plaster of Paris, gypsum is carefully heated. After losing 
three-fourths of its water of crystallization, the gypsum changes to 
plaster of Paris. 

2CaSO 4 2H 2 O < 3H 2 O + (CaSO 4 ) 2 H 2 O 


plaster of Paris 

Plaster of Paris, on the addition of water, combines with three mole- 
cules of water and re-forms gypsum. The reaction is thus reversible. 
This recombining of plaster of Paris with water is called setting, and 
is accompanied by the liberation of some heat. In preparing plaster 
of Paris care must be taken not to heat the gypsum above 175C., 
since all the water would then be driven off, and the gypsum would 
be said to be dead-burnt, and would be useless for plaster. 

Plaster of Paris is used in cheap statuary, casts for broken or dislo- 
cated bones, and molds of various objects such as teeth and metals. 
It is used also in plastering inside walls and ceilings, and in making 
plasterboards and similar partition materials. 

The superphosphate of calcium. In the discussion of the chief 
uses of chamber sulfuric acid, mention was made of the manufac- 
ture of calcium superphosphate. This compound was first prepared 




as a fertilizer from rock phosphate by John Lawes, in ltS42, and his 
patented process is still in use today. Calcium phosphate is insoluble 
in water, but the superphosphate is soluble and, hence, can be uti- 
lized by plants at once. 

Ca 3 (PO 4 )> 4- 2H 2 SO 4 - 2CaSO 4 + Ca(H ? PO 4 ) 2 


Superphosphate is made 1'roni rock phosphate which, within the 
United States, is mined chiefly in Florida, Tennessee, and South Car- 
olina. Florida rock phosphate (containing uranium impurities) is 
composed principally of Ca 3 (PO 4 ) 2 , the tricalcium salt of orthophos- 
phoric acid, H 3 PO 4 , a tribasic acid. Almost four million tons of rock 
phosphate are used in fertilizers each year in this country. Super- 
phosphate as shown in the equation above is really a mixture of 
monocalcium phosphate and gypsum, and contains the equivalent of 
about 17 percent P,,O r> . 

Triple superphosphate is a concentrated phosphate fertilizer de- 
veloped by the Tennessee Valley Authority (TV A) , a government 
agency created by the Congress in 1933. Since most soils in the Ten- 
nessee Valley are deficient in phosphorus, a plan was formulated for 
phosphate manufacture using the raw phosphate rock of Tennessee. 
Triple superphosphate is now widely used throughout the nation. It 
contains about three times as much available phosphate as ordinary 

Aerial view of a Florida 
phosphate rock plant. Under 
each of the storage piles are 
doors which, when opened, 
permit the rock to fall onto 
underground conveyor belts. 
The belts carry the rock into 
the plant for processing. 


UM by M 


superphosphate because the manufactured product contains no gyp- 
sum. It is made by the action of phosphoric acid on phosphate rock. 

Ca 3 (P0 4 ) 2 + 4H,P0 4 3Ca(H 2 P0 4 ) 2 

With this low-cost and highly-concentrated phosphate American 
farmers are winning the battle against an impoverished soil. 

Calcium phosphate as the source of pure phosphorus. Calcium 
phosphate is also a source of the pure phosphorus used in the manu- 
facture of phosphoric acid, P,() 5 , PC1 5 , and matches. The phosphorus 
is prepared by heating a mixture of calcium phosphate, sand, and 
carbon in a totally enclosed electric furnace. The calcium phosphate 
decomposes into calcium oxide, CaO, and phosphorus pentoxiclc, 
P,O 5 . The sand, SiO,, unites with the calcium oxide, forming cal- 
cium silicate slag, CaSi() a , which is drawn off from the bottom of the 
furnace. The carbon reduces the phosphorus pentoxide. forming 
phosphorus vapor and carbon monoxide. The phosphorus vapor is 
condensed under water and cast into sticks. The entire reaction may 
be represented as follows: 

Ca 3 (PO 4 ) 2 + 3SiO, + 5C - 3CaSiO 3 + 5CO + 2P 

Most of this phosphorus is burned in dry air to P 2 O 6 which is used 
to make phosphates including phosphoric acid. 

P,0 5 + 3H,0 - 2H : JP0 4 

This weak acid is used in soft drinks, jellies, and preserves. 

Hard water and calcium ions. Water containing dissolved calcium 
salts will not form a lather with soap until the calcium ions are 
removed. We call water containing such dissolved salts hard water. 
When the soluble calcium salts are precipitated as calcium soaps, the 




water will then form a lather with soap, and we say that we have 
softened the water. In other words, to soiten hard water, the free cal- 
cium ions must be removed. 

There are two kinds of hard water: (1) temporary hard water; 
(2) permanent hard water. Temporary hard water is water that con- 
tains calcium bicarbonate. Such water may he softened in two ways: 
(1) by boiling, and (2) by adding washing soda. When temporary 
hard water is boiled, the soluble calcium bicarbonate changes to in- 
soluble calcium carbonate, which precipitates out. Thus the free cal- 
cium ions are removed. The ionic equation is: 

Ca++ + 2(HCO 3 )- - CaCO 3 [ 4 HO + CO 2 1 

Washing soda softens temporary hard water in the following way: 
Ca-n- 4- 2(HCO a )- + 2Na+ + CO 8 ~ - CaCO 3 1 4 2Na+ + 2(HCO*)- 

Permanent hard water is water that contains calcium snlfate. Such 
water cannot be softened by boiling, but may be softened by the 
addition of washing soda, as the following ionic equation shows. 

.", - Ca++ 4 SO 4 ~ + 2Na+ + CO 3 -> CaCO 3 1 4- 2Na+ 4 SO 4 ~ 

Anotlier water-softener is Na,P() 4 (trisodium phosphate or TSP) , 
which forms insoluble calcium and magnesium phosphates, which 
are easily rinsed away. Glassy sodium metaphosphate, NaP() :j , which 
prevents the precipitation of calcium and magnesium soap, is found 
in dishwashing mixtures and is also used in municipal water supplies. 
The hexametaphosphate, (NaPO 3 ) , is marketed under the trade 
name, "Calgon." 

The Permutit Co, 

This composite photograph shows the 
difference between hair when washed 
in hard water (left) and washed in soft 
water (right). Note the soap scum de- 
posits in the photomicrograph (left). 



For many purposes, hardness of water need not be removed. For 
example, in drinking water the calcium ions are actually beneficial, 
since they supply an element necessary for bone-building. But in 
many industries using large quantities of water, the presence of cal- 
cium salts becomes a source of waste, and raises serious plant prob- 
lems. Calcium salts in boilers form deposits, or scales, which gradu- 
ally get thicker until they clog the pipes of the boiler. The efficiency 
of a boiler lined with a scale of calcium salts is greatly reduced and 
much heat is lost because the heat does not pass readily through the 
scale to heat the water. This may lead to explosions. 

If hard water is used in the home for laundering, bathing, wash- 
ing dishes, and so forth, the amount of soap required is much greater 
than if the water is soft. 

Although calcium salts are the most common cause of hard water, 
the salts of magnesium, iron, or barium also make water hard. What 
has been said about calcium salts holds true for the salts of these 
three elements also. Hardness in water is measured either in parts 
per million (ppm) or in degrees. One degree is equal to one grain 
of calcium salt per gallon of water. 

Ion-exchange method of softening water. One of the most efficient 
methods of softening water is by the use of "Permutit." This is the 
trade name for either a natural or synthetic compound composed of 
a complex sodium silico-aluminate called zeolite, which looks like a 
greenish-gray sand. Hard water is passed through a filter containing 
"Permutit"; the calcium ions replace the sodium ions, forming a 
calcium compound, calcium zeolite, which remains behind in the 
filter. There is no radical change in the structure of the solid zeolite 
during this interchange. 

Na 2 zeolite + CaSO 4 Ca zeolite + Na 2 SO 4 

Adapted from a drawing by The Permutit Company 

soft water out 
hard wafer in 

Fig. 106. Cross section of a zeo- 
lite home water-softener. 



When all the sodium in the "Permutit" has been replaced by cal- 
cium, the zeolite is covered with a ten percent solution of sodium 
chloride; the sodium of the sodium chloride replaces the calcium of 
the zeolite, which is now ready to soften more water. Again there is 
an ion exchange. 

Ca zeolite + 2NaCl -> CaCl 2 + Na 2 zeolite 

This interchangeability of Na and Ca ions enables "Permutit" to act 
as a water-softener for years without losing its effectiveness to any 
great extent. 

Demineralized water. Ion-exchange is also used in purifying water. 
Sodium zeolite removes calcium and magnesium ions but not the so- 
dium or chloride ions of salt or other ions such as sulfate and bi- 
carbonate that may be present. We now have other synthetic resin 
ion-exchangers that can rid water of these ions by a two-step process. 
Water is passed through alternate layers of two different exchangers 
and stripped of all its ionized salts. The first is a cation exchanger 
which removes the metallic ions and changes all salts present to their 
corresponding acids. In the following formula, the synthetic resin is 
represented by the symbol R. 

(1) HR + NaCl - NaR j + HC1 

The second is an anion exchanger that removes the free acids formed. 
In the following formula, the synthetic resin is represented by the 
symbol R'. 

(2) R' + HCl^R'-HCl j 

The water thus purified is called demineralized water and is almost 
as pure as distilled water. For household and many other uses it is 
just as good as distilled water. Ordinary tap water can be converted 
by means of a commercial demineralizer to almost pure water much 
more cheaply than by distillation. Even brackish water containing 
as much as 0.3 percent NaCl can be converted to water fit for irri- 
gation. Some day even sea water may be so treated economically. 

Synthetic resins are organic compounds that resemble natural 
resins such as amber. "Amberlites" is the trade name of one such 
group of resins. After use, they are regenerated and used over again. 
Ion-exchange is also used in the refining of sugar, and in the separa- 
tion and purification of vitamins, enzymes, and antibiotics such as 
streptomycin arid penicillin. It can even separate extremely minute 
amounts of amino acids in mixtures of these complex compounds 


Photomicrographs of soap dissolving in soft water (right) and trying to dissolve in 
hard water (left). Since soap is ineffective until dissolved, these photographs 
clearly illustrate why washing in hard water is unsatisfactory. 

by the recently developed technique called chromatography. For 
example, such a mixture which had hitherto defied separation is sent 
through a long glass column, well-packed with some adsorbent ma- 
terial such as AloO 3 , SiCX, or C. All ions are not held with equal 
strength by an ion-exchanger. The more strongly adsorbed amino 
acids are held at or near the top of the column, and the more weakly 
adsorbed, below in bands. Substances not adsorbed run through in 
the filtrate. Chromatography also has been used to separate some of 
the products of uranium fission, as well as the constituents of cell 


Huxley, Julian, and Andrade, E. N. daC. More Simple Sci- 
ence, pp. 154-222. Harper & Bros., New York, 1936. The chap- 
ters on "Soil" and "Agriculture" contain interesting and valu- 
able information on the general problem of soil enrichment 
as seen by two English scientists. 

Walton, Harold F. "Ion-Exchange." Scientific American, 
November, 1950, pp. 18-51. Explains the technique of re- 
placing electrically charged atoms or groups of atoms with 
others of the same charge. 

U.S. Department of Agriculture. Liming Soils for Better 
Fanning. Farmers' Bulletin No. 2032 (1951). Supt. of Docu- 
ments, U.S. Govt. Printing Office, Washington 25, D.C. 15^. 
A valuable illustrated pamphlet for the farmer. 




Group A 

1. Name the alkaline earth metals. 

2. Using diagrams showing the distribution of the elec- 
trons of Na and Ca, compare their chemical properties. 

3. (a) List the four forms of CaCO 3 . (b) Opposite each 
form write its chief use. 

4. (a) Under what conditions does insoluble CaCO 3 go 
into solution? (b) What salt is it that dissolves? Write the 
equation that shows how it forms, (c) When will the reaction 
begin to reverse? 

5. (a) Explain the formation of limestone caves, (b) How 
could a limestone cave be entirely filled with limestone by a 
natural process? 

6. (a) Describe a commercial method of preparing lime, 
(b) Write the equation. 

7. In the industrial preparation of quicklime, how is the 
reaction CaCO 3 - CaO + CO 2 made to go to completion? 


8. On what properties of lime was the use of limelight 

9. (a) On what physical properties does the use of quick- 
lime as the basic lining of various furnaces depend? (b) on 
what chemical property? (c) Write an equation that illustrates 
this chemical property. 

10. (a) What type of reaction is the slaking of lime? 
(b) Write the equation for it. 

11. (a) How is limewater made? (b) What are its prop- 
erties and uses? (c) What is whitewash? (d) What is milk of 

12. (a) Using an equation, explain how limewater is used 
as a test for CO . (b) What happens if you continue to pass 
CO 2 into limewater? (c) Write an equation explaining what 
happens, (d) How can this reaction be reversed? 

13. Ca appears below Na in the replacement series of the 
metals, (a) How would you expect Ca to react with water? 
(b) Compare its action on water with the action of Na on 

14. (a) If you left a barrel of lime outdoors on a rainy day, 
what would happen? (b) If the lime were then exposed to the 
air, what compound would be formed, at least in small quanti- 
ties? (c) Write equations showing the reaction in each case. 


15. Explain the old rhyme: 

"Lime and sand and water 
Make a very good mortar." 

16. Write two equations illustrating two changes that take 
place during the setting of mortar. 

17. Ca is an element essential in plant nutrition. Explain. 

18. How are calcium compounds used to "sweeten" soils? 

19. Explain the formation of the thin hard film on the sur- 
face of lime water in an open jar. 

20. What weight of lime can be produced from 200 tons of 
limestone containing 90 percent CaCO 3 ? 

21. CO 2 was passed through limewater and 150 g. of CaCO 3 
were precipitated. What volume of CO 2 was used? 

22. (a) How is bleaching powder made? (b) Write the 

23. Write an equation that shows both the manufacture of 
plaster of Paris and the setting of plaster of Paris. 

24. How would you tell crystals of gypsum from crystals of 

25. (a) What property of rock phosphate keeps it from be- 
ing used as a fertilizer? (b) How is it changed into a fertilizer? 
(c) Write the equation. 

26. Using equations, describe the manufacture of (a) P, 
and (b) H 3 PO 4 . 

27. The smell of C1 2 is detected when a fresh can of CaOCl 2 
is opened. Explain. 


28. Explain the "Permutit" process for softening water. 

29. (a) What do hard waters contain that makes them 
hard? (b) By what simple test can housewives tell whether 
water is hard? 

30. From the point of view of the dissociation theory, dis- 
tinguish between temporary and permanent hard water. 

31. Explain the two-step, ion-exchange process for demin- 
eralizing water. 

32. (a) How is temporary hard water softened in the home? 
Write the equation, (b) Which costs l*s: the softening of 
hard water or using large quantities of soap? 


33. (a) How is temporary hard water softened for indus- 
trial use? (b) Write the equation. 

Group B 

34. State three recent industrial applications of the ion- 
exchange process for removing undesirable ions. 

35. Write formulas of: (a) a calcium compound used as a 
dehydrating agent, and (b) a calcium compound used in the 
fixation of N 2 . (c) What combustible gas is made from this 
latter calcium compound? 

36. For the commercial bleaching of cotton cloth, a weak 
solution of sulfuric acid and bleaching powder is used. Chlor- 
ine is released. Write the equation for this reaction. 

37. (a) Describe fully a method you would use in the lab- 
oratory to prepare pure chalk, (b) What forces the reaction 
during which the chalk is formed to go to completion? 

38. Why do the walls of freshly-plastered houses "sweat"? 

39. What volume of CO 2 can be obtained from ten pounds 
of oyster shells containing 78% CaCO 3 ? 


1. Make an investigation of the method or methods used in 
your home or community to soften the water. Report to your 
class on your findings. If your community water supply is 
hard, make a survey of the cost of softening it. How would 
you justify such an expenditure? 

2. Bring a sample of water used for washing purposes in 
your home. Show the class whether it is hard or soft water, 
and, if hard, then soften it as a demonstration. Your teacher 
will supply you with the necessary material if you cannot ob- 
tain it yourself. 

3. Write a report on the present state and future possi- 
bilities of using sea water for irrigation and industrial pur- 
poses. (Read past issues of Chemistry published by Science 

4. Sprinkle some water over a few lumps of quicklime in a 
pan. Report your observations to the class. 





. . . In the "secretas" of t/ic twelfth 
century mention is made of sour 
galls, green vitriol (iron sulfate) , 
and a host of unimportant materials 
as being employed in the admixture 
of black inks. The first two (tanno- 
gallate of iron) , when used alone, 
form the sole base of all unadulter- 
ated gall inks. D. N. Carvalho, 1904 

Iron forms two series of compounds. As we have already learned, 
copper has two valences and forms both cuprous and cupric salts. 
Iron also has two valences. Ferrous ions carry two positive charges 
and ferric ions carry three positive charges. Chemists explain the 
existence of the ferrous and ferric ions in the same way that they 
explain the existence of cuprous and cupric ions. Because the distri- 
bution of the electrons differs in ferrous and ferric ions, iron forms 
two series of salts. Because the ferrous ion is less stable than the ferric 
ion, a ferrous salt tends to change into the corresponding ferric salt. 

How ferrous and ferric chlorides are related. A solution of ferrous 
chloride and also a solution of ferric chloride both contain chlorine 
ions. But the solutions differ in properties because of the presence of 
two different iron ions. A solution of ferrous chloride is almost color- 
less, while a solution of ferric chloride is brownish-yellow. There- 
fore, we may conclude that ferrous ions are almost colorless, while 
ferric ions are brownish-yellow. Ferrous chloride, FeCl 2 , is prepared 
by the action of iron on HC1 away from air. 

Fe + 2H+ + 2C1- 

2C1- + H 2 




Ferric chloride, used as an astringent in medicine, may be pre- 
pared by dissolving iron in hydrochloric acid and then passing chlo- 
rine through the resulting ferrous chloride solution. 

2Fe++ + 4C1- + C1 2 


From this equation we see that this reaction is an oxidation-reduction 
reaction. The iron in ferrous chloride has been oxidized to the ferric 
condition, since there has been an increase in valence from two to 
three (Fe f f to Fe+ ++ ). At the same time free chlorine with a valence 
of zero was changed to a valence of minus one. Therefore, chlorine 
has been reduced since it has gained an electron. 

Conversely, ferric chloride may be changed to ferrous chloride by 
the action of hydrogen, as shown by the following equation. 

Fe+++ + 3C1- + H -* Fe++ + 2C1~ + HC1 

Here the ferric ion has gained an electron and its valence has changed 
from three to two. Therefore, the iron in the ferric chloride has been 

Ferrous chloride, on exposure to the oxygen in air, changes gradu- 
ally into ferric chloride, the change being easily noted because of the 
change in color that accompanies the chemical reaction. 

How to identify ferrous and ferric ions. Further evidence that the 
ferrous ion differs chemically from the ferric ion is found in the dif- 
ferent tests that are used to identify each ion. A solution of potassium 
ferrocyanide, K 4 Fe (CN) 6 , added to a solution of a ferric salt forms 
ferric ferrocyanide, a deep blue precipitate called Prussian blue. 

4FeCl 3 + 3K 4 Fe(CN) 6 -> Fe 4 [Fe(CN) 6 ]3 


A solution of potassium ferricyanide, K 3 Fe (CN) , added to a fer- 
rous solution forms ferrous ferricyaiiide, a deep blue precipitate 

Glass rod 
Glass tube 

Slit in rubber tube 

Fig. 1O7. Laboratory prepara- 
tion of ferrous chloride. The 
narrow slit in the rubber tube 
allows the escape of hydro- 
gen from the test tube and 
prevents air from entering. 



whose composition and color, however, differ from prussian blue. 
This precipitate is called TurnbuWs blue. 

3FeCl 2 + 2K 3 Fe(CN) 6 Fe,[Fe(CN) 6 ] 2 1 + 6KC1 

These differing deep blue precipitates are specific tests tor ferric 
and ferrous ions. The test for a ferric salt will not work with a fer- 
rous salt, and vice versa. 

The reduction of ferric to ferrous salts is used in blueprints. 
The reversible change of ferrous to ferric compounds is made use 
of in the manufacture of blueprint paper and ink. Architects and 
engineers use blueprints instead of the original plans. The process of 
making blueprints was invented by Sir John Herschel, the eminent 
English astronomer. White paper is coated with a thin film of (1) fer- 
ric ammonium citrate, and (2) potassium ferricyanidc. This sensi- 
tized paper is then covered with tracing paper containing the design 
to be reproduced. In darkness, reaction does not occur because a fer- 
ric salt does not react with potassium ferricyanide in the absence of 

When exposed to light, the ferric ammonium citrate is reduced to a 
ferrous salt, which then reacts with the potassium ferricyanide, 
K 3 Fe(CN),,, forming deep TurnbuH's blue. In places where the 
sensitized paper has been exposed to light, that is, wherever the 
design on the tracing paper has not protected it, the paper will be 
blue. In places where light has not reached it, the paper remains 
white. After exposure, the excess chemicals are washed off in running 
water, leaving the blue precipitate clinging to the paper. 

The oxidation of ferrous to ferric salts is used in making writing 
ink. The commonly used blue-black writing ink is a water solution 

Ook rich In tcm- 

nic are with 

mm4 *o 

Wmtwrupli*, The Carter'* Ink *>/.- 



of an iron salt of tannic acid held in suspension by a gummy sub- 
stance, or colloid, called gum arable. A small amount of carbolic 
acid is added to prevent the growth of molds. Such an ink is prepared 
by adding green ferrous sulfate called copperas, FeSO 4 7H 2 O, to 
tannic acid obtained from nutgalls, a kind of abnormal growth on 
trees caused by the sting of certain insects such as wasps. The ferrous 
tannate that is formed is almost colorless. On exposure to air, how- 
ever, the ferrous tannate changes to ferric tannate, which is a black 
insoluble substance. Since the black ferric tannate is not formed un- 
til the ink has had a chance to oxidize on the paper, a temporary 
blue coloring matter, such as indigo, is added to the ink. The color 
change, therefore, is from blue to black. The chemical equation may 
be represented thus: 

Ferrous tannate + oxygen (from air) > ferric tannate 
colorless black 

Removing stains made by ferro-tannate inks. Ferro-tannate inks 
are permanent inks, that is, they are not affected by oxygen or water. 
To remove stains made with this kind of ink, it is necessary to reduce 
the black ferric tannate with some suitable reducing agent, such as 
oxalic acid (oxalates are present in rhubarb) , lemon juice (citric 
acid) , or sour milk (lactic acid) . The soluble ferrous tannate and 
the other ingredients of the ink may then be washed away. 

In using oxalic acid to remove a stain made by a permanent ink, 
soak the stain for a few seconds in a solution of three tablespoonfuls 
of oxalic acid crystals dissolved in one pint of water. Rinse by dip- 
ping first in hot water and then in water to which a few drops of 
household ammonia have been added. This method of stain removal 
should not be used on weighted silk fabrics, since oxalic acid attacks 
silk fibers. Solutions of bleaching powder or of sodium hypochlorite 
also are used effectively in removing the stains made by permanent 

Some inks do not contain iron salts. Not all inks contain iron salts. 
For example, laundry ink used for marking linen consists of an am- 
moniacal solution of silver nitrate mixed with a small amount of 
provisional, or temporary, color and a gum for thickening the ink. 
On exposure to light, the silver nitrate eventually changes to a very 
finely divided colloidal silver, which is black and is unaffected by 
ordinary reagents. India ink, indelible ink, and printing ink all con- 
sist of carbon (lampblack) held in colloidal suspension by gum ara- 
bic or some other colloid. They are permanent, since neither acids 
nor bases dissolve carbon. 

U.S.D.A. Bureau of Home. Nutrition and Home Economics 

Javelle water is used to remove certain kinds of stains. 

Removing stains made by printing ink. To remove a stain made 
by printing ink, rub the stain with lard or Vaseline and work it well 
into the stain. This softens the medium in which the carbon pigment 
is held and, to some extent, disperses the carbon. If the stained ma- 
terial is washable, wash well with soap and water. Otherwise, sponge 
with carbon tetrachloride, ethylene trichloride,. or gasoline. 

Kerosene alone will frequently remove printing ink stains. Soak 
the stained material in kerosene for several hours, wash thoroughly 
with soap and hot water, rinse, and dry. Since this method of stain 
removal is rather vigorous, it is recommended for rugged fabrics only. 

Washable inks. Some inks are solutions of certain aniline dyes, 
such as cosine (red) or nigrosine (black) crystals. As such inks are 
usually temporary, or washable, inks, they are not suitable for mak- 
ing permanent records. They fade after long exposure to air and 
stains made by them may be washed away with water. 

Sympathetic inks. A concentrated water solution of cobalt chloride 
is pink (color of the cobalt ion) , yet the anhydrous salt is blue. 
Therefore, paper on which a dilute solution of this salt has been 
used for writing will show a very pale pink, practically colorless 
writing when damp, and a clear blue writing when the paper is dried 
by being warmed. Hence it can be used as a sympathetic ink; that 
is, it may be changed from invisible to visible by proper treatment. 

Pellets of alumina impregnated with CoCL are used in some salt 
shaker caps. The CoCl 2 changes from blue to pale pink as it absorbs 
moisture. It can be heated in an oven to drive out the water, and 
used over again. 

Another kind of sympathetic ink was used during the Indian Mu- 
tiny in 1857. The besieged English soldiers sent out secret messages 
through the Indian lines by writing with colorless rice water. The 



receiver of the message dipped the paper in weak iodine solution 
and writing appeared. This effect was caused by the precipitation of 
a bluish-black substance by the reaction of the rice starch and iodine. 

Laundry bluing. Most white fabrics have a slightly yellowish tint. 
The use of bluing does not remove this tint but in combination with 
the yellow already present, merely masks it by producing a taint 
gray that appears white. One of the many kinds of soluble bluings 
in use consists of Prussian blue. Others are blue dyes. 

Life too, depends on this reversible change of Fe +++ and Fe++ . 
In almost all living cells respiration is controlled by a change 
of Fe 4 ++ to Fe +f containing catalysts called cytochromes. The 
Fe ++ + picks up an electron from, let us say, sugar and changes to the 
Fe++ condition. The Fe+ + cytochrome transfers this electron to oxy- 
gen thus returning to the Fe+++ state. 

Lodestone, and two other oxides of iron. A most interesting ore 
is lodestone, an oxide of iron whose formula is Fe ;! () 4 . The ancients 
knew about the valuable magnetic property of this iron compound, 
and among their writings we find: "It not only attracts iron rings 
but also imparts to them a similar power of attracting other rings, 
and sometimes you may see a number of pieces of iron suspended 
from one another so as to form a long chain." The magnetic prop- 
erty of lodestone was first used by mariners in the magnetic compass, 
which enabled them to venture far out to sea. It may be made by 
passing steam over hot iron. 

3Fe + 4H 2 O - Fe 3 O 4 + 4H 2 1 

Ferric oxide, Fe 2 O 3 , occurs abundantly as hematite. It may also 
be prepared by roasting a ferrous compound. Ferric oxide is used 

International Print inn Ink 

The first step in making printing ink is 
mix a chemical pigment with a vehicle such 
as linseed oil. 


as a red pigment in paints, and as jeweler's rouge for polishing 
metals. It is frequently the red pigment in rouge, lipstick, and other 
cosmetics. Iron rust is the hydra ted form of hematite. Its formula 
may be written (Fe 2 O 3 ) 2 . 3H 2 O. 

The brown or red color of certain soils, particularly in Georgia 
and certain sections of Minnesota, is caused by iron compounds. 
Most of the iron made in America before and during the Revolu- 
tionary War was produced by the reduction of the relatively low 
grade hematite ores on the Atlantic seaboard. 

Ferrous oxide, FeO, a black powder which readily oxidizes to fer- 
ric oxide, may be obtained by reducing ferric oxide with hydrogen. 
The equation is: 

Fe 2 3 + H 2 - 2FeO + H 2 O 


National Bureau of Standards. Inks. Circular No. 95. Wash- 
ington, D.C. 

U.S. Department of Agriculture. Stain Removal from Fab- 
rics: Home Methods. Catalogue No. A1.9: 1474/3. Supt. of Doc- 
uments, Washington, D.C. l(ty. Discusses the removal not only 
of ink stains but of many other common stains. 


Group A 

1. In what ways does the ferrous ion differ from the ferric 
ion, physically and chemically? 

2. Using equations, tell how ferrous chloride can be pre- 
pared (a) from iron, (b) from ferric chloride. 

3. Using an equation, show how ferric chloride may be 
prepared from ferrous chloride. 

4. What is the relation between change of valence and 

5. What is the relation between change of valence and re- 

6. According to the electron theory, define and illustrate 

7. According to the electron theory> define and illustrate 


8. When Zn dissolves in H 2 SO 4 , is the Zn oxidized or re- 
duced? Explain. 

9. A solution contains an iron salt. How would you de- 
termine whether the salt is a ferrous or a ferric compound? 

10. What is the valence of the first iron in Prussian blue, 
Fe 4 [Fe(CN) 6 ],? 

11. What is the valence of the ferricyanide ion in Turn- 
bull's blue, Fe ? [Fe (CN) ] 2 ? 

12. A solution contains both ferrous salts and ferric salts. 
How would you test for the ferrous salt? 

13. What uses are made of the reduction and oxidation of 
iron salts? 

14. (a) What is the light-sensitive chemical with which blue- 
print paper is covered? (b) Light changes it into what other 

15. Explain how blueprinting is done. 

16. Why is it necessary to wash blueprint paper after ex- 

17. What is the chief chemical change that takes place in 
blue-black writing inks? 

18. Name the functions of gum arabic and carbolic acid 
in writing ink. 

19. (a) How would you remove fresh ink stains from cotton 
cloth? (b) What type of reaction is involved? 

20. What are black ink crystals? 

21. Describe one kind of sympathetic ink. 

22. (a) What is the composition of laundry ink? (b) What 
chemical change does it undergo during drying? 

23. What three inks contain lampblack as a base? 

24. (a) Can lampblack ink stains be removed by reducing 
agents? (b) Why? 

25. What are the formulas and names of three iron oxides? 

26. Name two washable inks. 

27. What is the relation between the change of Fe++ to 
Fe+++ and respiration? 


28. Write the equation for the reaction between hot iron 
and steam. 

29. What is the composition of (a) jeweler's rouge, 
(b) lodestone, (c) rust? 

30. Why are ferric compounds more stable than ferrous com- 

31. What two oxidizing agents can be used to change a fer- 
rous salt to a ferric salt? 

32. Determine the percentag